COLLEGE  TEXTBOOK 

OF 

CHEMISTRY 


BY 

WILLIAM  A.  NOYE£> 


DIRECTOR  OF  THE   CHEMICAL  yA 

OF  THE   UNIVERSITY    OF  ILLINOIS 


NEW  YORK 
HENRY  HOLT  AND   COMPANY 


COPYRIGHT,  1919 

BY 
HENRY  HOLT  AND  COMPANY 


PREFACE 

This  College  Textbook  of  Chemistry  is  designed,  more  es- 
pecially, for  students  of  the  Freshman  or  Sophomore  years 
in  college  who  have  not  studied  chemistry  in  the  High 
School.  As  with  all  textbooks  for  beginners,  two  purposes- 
have  been  constantly  kept  in  mind  while  writing  the  book: 
the  presentation  of  a  few  of  the  multitude  of  chemical  facts 
which  touch  our  modern  life,  in  such  a  manner  that  they 
can  be  clearly  understood,  and  the  discussion  of  the  theories 
and  principles  around  which  all  our  chemical  knowledge  is 
grouped.  The  teacher  of  chemistry  is  embarrassed  by  the 
vast  and  ever  increasing  amount  of  knowledge  at  his  dis- 
posal and  is  often  tempted  to  present  many  more  topics  than 
the  student  can  possibly  remember.  In  trying  to  avoid 
this  difficulty  many  facts  ordinarily  included  in  an  elemen- 
tary textbook  have  been  omitted  and  those  which  are  given 
are  brought  as  far  as  possible  into  close  logical  relations. 

The  summary  at  the  close  of  each  chapter  is  a  somewhat 
unusual  feature  of  the  book.  It  is  hoped  that  these  summa- 
ries will  be  found  useful. 

Success  in  the  study  of  chemistry  depends  especially  on 
the  ability  to  learn  new  facts  in  their  relation  to  those  which 
have  already  been  acquired  and  on  the  cultivation  of  a 
logical  as  distinguished  from  an  arbitrary  memory.  The 
exercises  at  the  close  of  each  chapter  and  questions  occa- 
sionally inserted  in  the  text  are  designed  to  assist  the  student 
in  this  direction.  These,  and  similar  exercises  prepared  by 
the  teacher,  should  be  given  very  careful  attention  in  the 
class  room. 


435145 


vi  PREFACE 

I  wish  to  express  my  very  sincere  thanks  to  Dr.  B.  S. 
Hopkins,  Dr.  H.  G.  Deming,  Dr.  Charles  Davidson  and 
Mrs.  H.  A.  Davidson,  who  have  read  the  manuscript  and 
made  many  valuable  and  helpful  criticisms;  also  to  Dr.  J.  H. 
Ready  and  Dr.  B.  S.  Hopkins,  who  have  read  the  proofs, 
and  to  Dr.  Charles  Davidson,  who  has  prepared  the  index. 


CONTENTS 

CHAPTER  PAGE 

I.  Fundamental  Principles 1 

II.  Combustion 6 

III.  Hydrogen     . 19 

[V.  Weights  and   Measures,   Molecular  Theory,  Laws  of 

Gases 28 

Composition   of    Water.     Laws    of    Composition   by 

Weight.     The  Atomic  Theory 42 

VI.  Properties  and  Uses  of  Water.     Vapor  Pressure  ...      58 

VII.  Sodium,  Acids,  Bases,  Salts 72 

VIII.  Hydrochloric  Acid,  Chlorine,  Oxygen  Acids  of  Chlorine     80 
IX.  Group  VII;  the    Halogen    Family,  Bromine,  Iodine, 

and  Fluorine 92 

X.  Classification  of  Elements.     Valence 100 

XI.  Group  VI;  Sulfur,  Selenium  and  Tellurium 109 

XII.  Selection  of  Molecular  and  Atomic  Weights      ....    130 

XIII.  Group  V;  Nitrogen 138 

XIV.  Air.     The  Noble  Gases.     Group  Zero    " 155 

XV.  The  Periodic  System      162 

XVI.  Group  V;  Phosphorus,  Arsenic,  Antimony  and  Bismuth  170 

XVII.  Group  IV;  Carbon 185 

XVIII.  Hydrocarbons,  Gas,  Flame 195 

XIX.  Carbon  Monoxide,  Carbon  Dioxide,  Carbon  Disulfid, 

Cyanides 211 

XX.  Carbohydrates,  Alcohols,  Acids,  Bread,   Proteins,  Di- 
gestion, Antitoxins,  Alkaloids,  Dyes 219 

XXI.  Group  IV;  Silicon,  Tin  and  Lead 237 

XXII.  Group  III;  Boron,  Aluminium 253 

XXIII.  Group  II;  First  Division;  Alkali   Earth-metals,  Cal- 
cium, Strontium,  Barium,  Radium 260 

XXIV.  Metallurgy  and  the  Preparation  of  Compounds  of  the 
Metals 277 

XXV.  Group  II;  Second  Division;  Magnesium,  Zinc,  Cad- 
mium and  Mercury 283 

XXVI.  Group  I;  First  Division;  Alkali  Metals:  Sodium,  Potas- 
sium and  Ammonium.     Spectrum  Analysis  .    .    .    :    .   290 


viii  CONTENTS 

^  CHAPTER  PAGE 

XXVII.  Group  I;  Second  Division;  Copper,  Silver,  Gold;  Pho- 
tography  308 

XXVIII.  Group   VI;   Second   Division;   Chromium,    Tungsten, 

Uranium  ...... 316 

XXIX.  Group  VII;  Second  Division;  Manganese 320 

XXX.  Group  VIII;  Iron,  Cobalt,  Nickel,  Platinum     ....   323 
XXXI.  Analysis .335 


COLLEGE  TEXTBOOK  OF  CHEMISTRY 


CHAPTER  I 
FUNDAMENTAL  PRINCIPLES 

Science  is  systematic  knowledge.  Especially  it  is  that 
sort  of  knowledge  which  discovers  relations  between  the 
facts  or  phenomena  of  experience  and  expresses  these  rela- 
tions in  the  form  of  laws. 

Gravitation.  Laws. — It  is  a  common  observation  that 
an  object  which  is  not  supported  will  fall  toward  the  earth. 
Newton,  one  of  the  greatest  scientific  men  of  jthe  world, 
observed  an  apple  fall  to  the  ground.  By  reflecting  on  this 
observation  and  by  stud}dng  the  rates  of  the  motions  of  the 
planets  about  the  sun  he  discovered  the  law  of  gravitation — 
that  all  material  objects  attract  each  other  with  a  force 
which  is  proportional  to  their  masses  and  inversely  pro- 
portional to  the  squares  of  their  distances  apart.  Laws 
which  have  been  discovered  in  a  similar  manner  by  ob- 
servation, experiment  and  reflection  furnish  the  nucleus 
about  which  all  our  scientific  knowledge  is  grouped  in  a 
systematic  manner. 

Physical  Science  is  that  part  of  our  systematic  knowledge 
which  deals  with  the  phenomena  of  inanimate  nature. 
These  phenomena  may  be  considered  from  two  quite  differ- 
ent points  of  view. 

(a)  Physics. — We  may  direct  our  attention  to  the  motion 
of  bodies  and  to  the  phenomena  of  sound,  heat,  light  and 
electricity,  or,  in  general,  toward  those  phenomena  which 


2  ;  .:  t .  FUNDAMENTAL  PRINCIPLES 

we  group  together  under  the  general  name,  energy.  In  its 
simplest  form  energy  is  merely  the  motion  of  bodies  or 
some  force  which  may  produce  motion.  The  motion  of  a 
body  is  called  kinetic  energy.  A  force  which  may  produce 
motion,  directly  or  indirectly,  is  called  potential  energy. 
There  may  be  some  question  whether  there  is  any  sort  of 
energy  which  is  not  of  the  nature  of  motion.  The  branch 
of  science  which  considers  the  subject  of  energy  and  its 
transformations  is  called  physics. 

(6)  Chemistry. — From  a  quite  different  point  of  view  we 
may  consider  the  materials  of  which  substances  are  com- 
posed, distinguishing  pure,  or  homogeneous,  substances  from 
mixtures,  and  studying  how  such  substances  act  on  each 
other  under  various  conditions.  In  other  words,  we  may 
direct  our  attention  to  the  composition  of  substances.  The 
branch  of  science  which  does  this  is  called  chemistry. 

Mixtures. — The  first  step  in  the  study  of  substances  in  a 
chemical  way  is  the  separation  of  pure  substances  from  the 
mixtures  which  are  met  in  common  experience.  Two  of 
the  most  effective  means  used  for  this  purpose  are  crystalli- 
zation and  distillation. 

The  salt  brines  found  at  Syracuse,  New  York,  in  Michigan 
and  in  other  parts  of  the  world  contain  a  number  of  other 
substances  in  solution  with  the  salt,  but  when  these  brines 
are  evaporated  very  nearly  pure  crystals  of  salt  separate, 
while  the  other  substances  remain  in  solution.  On  the 
other,  hand,  if  the  steam  which  is  boiled  away  from  the 
brine  is  condensed,  water  free  from  salt  can  be  obtained. 
If  a  mixture  of  alcohol  and  water  is  boiled,  the  portion  of  the 
mixture  which  distils  first  will  contain  more  alcohol  than 
the  portion  remaining  behind  and  by  repeating  the  distilla- 
tion several  times  it  is  possible  to  separate  the  original 
mixture  into  nearly  pure  alcohol  and  pure  water. 

Pure  Substances  are  Homogeneous,  that  is,  all  parts  of 
them  are  alike.  If  they  melt  or  boil  without  decomposition 


LAW  OF  CONSTANT  PROPORTION         3 

each  has  a  definite  melting  point  or  boiling  point,  under 
atmospheric  pressure.  They  always  act  in  the  same  manner 
when  brought  in  contact  with  a  given  other  substance 
under  the  same  conditions.  If  they  are  compounds  they 
always  have  the  same  composition  by  weight. 

Law  of  Constant  Proportion. — The  last  property  of  pure 
substances  mentioned  has  been  established  by  the  accurate 
analysis  of  many  different  compounds  and  these  analyses 
have  established  one  of  the  most  important  laws  of  chemis- 
try— the  law  of  constant  proportion.1 

Compounds.  Synthesis. — If  fine  iron  filings  and  sulfur 
are  thoroughly  mixed  in  a  mortar  a  mixture  is  obtained 
which  has  a  uniform  gray  color  and  which  appears  to  the  eye 
homogeneous.  With  a  microscope,  however,  it  is  easy  to 
see  that  the  substance  is  still  a  mixture  of  particles  of  iron 
with  particles  of  sulfur.  With  a  magnet  the  particles 
of  iron  may  be  picked  up  and  the  particles  of  sulfur  which 
still  cling  to  the  iron  may  be  blown  away.  If  the  mixture 
is  heated  over  a  flame,  after  a  short  time  it  will  begin  to 
glow  and  the  mass  will  become  red  hot.  After  cooling,  the 
material  is  homogeneous  and  neither  particles  of  iron  nor 
particles  of  sulfur  can  be  seen  with  a  microscope.  A  mag- 
net will  no  longer  attract  the  iron  which  the  substance 
contains.  The  new  substance  is  evidently  formed  by  the 
union  of  iron  and  sulfur.  It  is  called  iron  sulfide  and  is 
spoken  of  by  chemists  as  a  compound  of  iron  and  sulfur. 
Such  a  process  of  putting  two  substances  together  and 
making  from  them  a  different,  new  substance  is  called 
synthesis. 

Direct  Analysis. — If  oxide  of  mercury  is  heated  in  a  test- 
tube  it  will  be  decomposed  into  metallic  mercury,  which 
will  condense  on  the  walls  of  the  tube,  and  a  gas,  oxygen, 
which  will  cause  a  glowing  splinter  to  burst  into  flame. 
By  this  process  the  compound,  oxide  of  mercury,  is  decom- 

1  Compounds  of  isotopes  are  an  exception  to  this  law,  see  p.  273.- 


4  FUNDAMENTAL  PRINCIPLES 

posed  into  two  new  substances,  mercury  and  oxygen,  and 
the  process  may  be  called  a  direct  analysis  of  the  oxide  of 
mercury.  The  object  of  any  chemical  analysis  is,  of  course, 
to  determine  the  substances  which  enter  into  the  com- 
position of  some  compound  or  mixture. 

Elements.  Compounds. — Nearly  all  of  the  substances 
which  we  find  in  the  world,  or  which  may  be  prepared  by 
artificial  means — more  than  one  hundred  thousand  are 
known  to  exist — may  be  separated  into  two  or  more  different 
substances,  just  as  oxide  of  mercury  may  be  separated  into 
mercury  and  oxygen,  though  in  most  cases  the  methods 
required  for  their  separation  are  not  so  simple.  Many 
substances  may  be  prepared  by  synthesis  as  sulfide  of  iron 
is  prepared  from  sulfur  and  iron.  But  a  comparatively 
small  number  of  substances  have  been  found  which  no  one 
has  been  able  to  separate  into  two  or  more  other  substances 
and  which  no  one  has  been  able  to  prepare  by  putting  two 
or  more  other  substances  together.  Such  substances  are 
called  elements. 

The  elements  which  have  been  longest  known  and  which 
are  most  familiar  in  our  daily  experience  are  the  metals 
— iron,  copper,  tin,  zinc,  lead,  mercury,  silver  and  gold. 
Another  set  of  elements,  which  were  not  known  till  compara- 
tively modern  times,  includes  gases,  oxygen  and  nitrogen 
of  the  air,  hydrogen,  obtained  by  decomposing  water,  and 
chlorine,  a  constituent  of  salt.  Sulfur,  which  is  used  in 
making  sulfuric  acid,  or  oil  of  vitriol,  phosphorus,  used  in 
matches,  and  carbon,  which  is  nearly  pure  in  charcoal, 
graphite  and  diamonds,  are  also  known  to  most  persons 
before  they  begin  the  systematic  study  of  chemistry.  If  we 
add  to  these  silicon,  found  in  sand,  aluminium,  an  element 
found  in  clay,  calcium,  found  in  limestones,  magnesium, 
also  found  in  many  limestones,  sodium  from  salt,  and  potas- 
sium from  wood  ashes,  we  have  a  list  of  twenty-one  elements. 
About  ninety  eight  per  cent  of  the  part  of  the  earth  which 


CHEMICAL  ELEMENTS  5 

we  can  examine  and  of  the  air  above  is  composed  of  these 
elements,  and  very  few  others  require  more  than  a  passing 
mention  in  an  elementary  textbook. 

About  sixty  other  elements  have  been  discovered  and 
we  now  have  strong  reasons  for  believing  that  the  number 
-of  elements  in  the  universe  does  not  much  exceed  one 
hundred. 

SUMMARY 

Science  is  systematic  knowledge. 

A  law  is  a  general  relation  between  facts  or  phenomena. 

Physics  deals  primarily  with  energy  and  its  transfer. 

Chemistry  considers  the  composition  of  substances. 

Pure  substances  are  most  often  obtained  by  crystalliza- 
tion or  distillation. 

Pure  substances  are  homogeneous,  have  a  constant  melt- 
ing point  or  boiling  point,  always  act  in  the  same  way  under 
the  same  conditions,  and  have  a  constant  composition. 

Compounds  are  substances  which  can  be  separated  into 
two  or  more  other  substances. 

Elements  cannot  be  so  separated. 

The  composition  of  a  compound  may  be  determined  by 
synthesis,  by  direct  analysis,'  or  by  indirect  analysis. 

Fewer  than  one  hundred  elements  are  known  and  of  these 
about  twenty  are  of  first  importance. 

QUESTIONS 

1.  What  other  substances  beside  salt  are  purified  by  crystal- 
lization? 

2.  What  substances  beside  water  and  alcohol  are  separated  by 
distillation? 

3.  What   elements    are   contained   in   the   following   common 
substances:  starch,  coal,  brass,  solder,  kerosene,  albumin,  §and, 
snow,  air? 

4.  Which  is  the  most  abundant  element?     Which  is  second? 
Which  is  most  important  for  life? 


CHAPTER  II 
COMBUSTION 

Fire. — It  is  a  common  experience  to  see  wood  or  coal 
burn,  leaving  ashes  that  are  very  much  lighter  than  the 
material  which  was  burned.  The  oil  of  lamps  and  the 
paraffin  or  tallow  of  candles  seem  to  disappear  as  they 
burn,  leaving  no  visible  products.  These  phenomena  have 
been  familiar  during  the  hundreds  of  thousands  of  years 
which  have  passed  since  the  human  race  first  learned  the 
use  of  fire,  but  it  was  less  than  a  century  and  a  half  ago  that 
a  satisfactory  explanation  of  burning,  which  we  call  combus- 
tion, was  discovered— so  very  recent,  in  comparison  with 
the  age  of  the  race,  is  our  knowledge  of  some  of  the  most 
fundamental  facts  of  chemistry. 

Relation  of  Air  to  Combustion. — It  was  finally  shown  that 
the  disappearance  of  the  wood  or  oil  is  not  due  to  the  de- 
struction of  the  matter  which  they  contain,  but  that,  instead 
of  this,  invisible  substances  are  formed  which  weigh  more 
than  those  burned.  Most  people  have  noticed  the  film  of 
moisture  which  is  formed  on  a  cold  lamp  chimney  when  the 
lamp  is  first  lighted.  This  is  due  to  water  which  is  formed 
by  the  burning  oil.  By  drawing  the  air  from  above  a  burn- 
ing lamp  through  clear  lime  water  it  can  easily  be  shown 
that  another  substance,  carbon  dioxide,  is  also  formed. 
If  these  two  substances  are  absorbed  by  soda  lime  and 
weighed,  as  may  be  done  with  the  apparatus  shown  in 
the  figure  (Fig.  1),  it  will  be  found  that  they  weigh  more 
than  the  oil  or  the  part  of  the  candle  which  has  been  burned. 
A  further  study  of  the  matter  has  shown  that  the  burning 

6 


LAVOISIER'S  EXPERIMENT  7 

oil  or  candle  takes  up  oxygen  from  the  air  and  that  if  all 
of  the  water  and  carbon  dioxide  are  collected  they  weigh 
exactly  as  much  more  than  the  oil  or  candle  as  the  weight 
of  the  oxygen  which  has  been  taken  up.  The  substances 
which  disappear  in  burning  change  their  form  but  there  is 
no  change  in  their  weight. 


FIG.  1. 

Lavoisier's  Experiment. — In  1775  the  French  chemist, 
Lavoisier,  showed  clearly  for  the  first  time  the  relation 
between  air  and  combustion.  Shortly  before  that  Priestly, 
an  Englishman,  had  prepared  oxygen  gas  by  heating  a  red 
compound,  which  we  now  call  oxide  of  mercury.  He  had 
shown  that  substances  which  burn  in  air  burn  with  much 
greater  brilliancy  in  oyxgen  gas.  Lavoisier,  after  hearing 
from  Priestly  of  this  experiment,  devised  the  apparatus 
shown  in  the  figure.  He  placed  some  mercury  in  the  retort, 
A,  and  bent  the  wide  neck  of  the  retort  in  such  a  manner 
that  the  bell-jar,  FG,  might  have  its  rim  below  the  surface 
of  mercury  in  the  dish,  S,  while  air  could  still  pass  freely 


8 


COMBUSTION 


from  the  interior  of  the  retort  to  the  interior  of  the  bell-jar. 
He  then  heated  the  mercury  in  the  retort  by  means  of  the 
charcoal  furnace,  MN,  for  several  weeks.  Red  oxide  of 
mercury  gradually  accumulated  on  the  surface  of  the 
mercury  in  the  retort,  while  the  mercury  rose  in  the  bell-jar, 
showing  that  a  part  of  the  air  was  disappearing.  After 
some  weeks  the  mercury  rose  no  further  on  continued 
heating  and  the  experiment  was  stopped.  Lavoisier  then 
noticed  carefully  how  far  the  mercury  had  risen  and  how 
much  of  the  air  had  been  absorbed.  When  the  oxide  of 


M 


FIG.  2. 

mercury,  which  he  collected^carefully,  was  heated  till  it  was 
all  decomposed,  he  found  that  he  had  obtained  a  volume  of 
oxygen  gas  equal  to  the  decrease  in  volume  in  the  air  of  the 
retort  and  bell-jar.  In  this  way  it  was  shown  that  about 
one-fifth  of  the  volume  of  the  air  consists  of  oxygen  and 
four-fifths  of  some  other  gas.  which  does  not  act  on  the 
mercury  or  support  combustion.  This  other  gas  is  chiefly 
nitrogen. 

Preparation  of  Oxygen. — The  preparation  of  oxygen  by 
heating  oxide  of  mercury  has  been  referred  to.  This  method 
of  getting  the  gas  is  tedious  and  expensive  and  it  is  only 


CATALYSIS 


9 


used  on  a  small  scale  in  the  laboratory  or  lecture  room, 
because  of  its  simplicity  and  because  of  the  historical 
interest.  For  ordinary  laboratory  uses,  small  amounts 
of  oxygen  are  prepared  by  heating  potassium  chlorate, 
a  white,  crystalline  salt,  which  melts  rather  easily  when 
heated.  It  decomposes  into  potassium  chloride  and  oxygen, 

slowly  at  its  melting  point,  more  rapidly  at      ^ 

higher  temperatures.  If  some  manganese 
dioxide  is  mixed  with  the  potassium  chlorate, 
the  decomposition  is  very  much  hastened, 
though  the  manganese  dioxide  is  not,  itself, 
permanently  changed.  To  prepare  oxygen 
in  the  laboratory  a  mixture  of  potassium 
chlorate  and  manganese  dioxide  is  heated 
in  a  test-tube,  or  in  a  glass  retort  or  flask, 
or  in  a  copper  flask. 

Catalysis. — Manganese  dioxide  hastens 
the  decomposition  of  potassium  chlorate 
but  is  not  permanently  changed  in  the 
process.  The  action  can  be  easily  shown 
by  melting  some  potassium  chlorate  in  a 
test-tube  and  dropping  in  some  of  the  pow- 
dered dioxide.  A  vigorous  decomposition 
of  the  chlorate  will  set  in  at  once.  A 
substance  which  hastens  a  chemical  action  by  its  presence, 
in  this  manner,  is  called  a  catalytic  agent  or  a  catalyst. 

Collection  and  Storage  of  Gases. — Gases  like  oxygen, 
which  are  only  slightly  soluble  in  water,  are  collected  in 
bottles  or  in  a  gasometer  (Fig.  3).  The  gasometer  is  filled 
with  water  by  opening  the  stopcocks  A,  allowing  the  water 
to  run  to  the  bottom,  and  B,  which  allows  the  air  in  the 
gasometer  to  escape.  When  the  gasometer  is  full  of  water 
the  two  stopcocks  are  closed,  the  cap  C  is  removed  and  the 
tube  which  delivers  the  oxygen  is  inserted.  As  the  opening 
closed  by  the  cap  slants  upward,  air  cannot  enter  but 


FIG.  3. 


10  COMBUSTION 

the  water  escapes  through  it  as  the  water  in  the  upper  part 
is  displaced  by  the  gas.  After  filling,  the  cap  is  replaced 
and  the  stopcock  A  is  opened.  The  gas  may  then  be  drawn 
off  through  B. 

Oxygen  from  Liquid  Air. — As  shown  by  Lavoisier's 
experiment,  air  is  composed  mainly  of  oxygen  and  nitrogen, 
about  one-fifth  of  its  volume  being  oxygen.  When  air 
which  has  been  compressed  is  allowed  to  expand,  it  grows 
cold.  Linde  and  others  have  devised  machines  by  means  of 
which  air  can  be  compressed  to  a  pressure  of  2000  to  3000 
pounds  to  the  square  inch.  The  air  is  then  allowed  to 
expand  through  a  small  copper  tube  several  hundred  feet 
in  length.  This  is  so  arranged  that  the  cold  air  escaping 
from  the  lower  end  passes  back  over  the  outside  of  the 
tube  and  thus  cools  still  more  the  current  of  air  passing 
down  the  tube.  In  this  way  the  air  soon  becomes  so  cold 
that  part  of  it  becomes  liquid.  When  this  liquid  air, 
which  is  at  a  temperature  of  about  185° C.  (or  300°  F.)  below 
zero,  is  allowed  to  boil,  the  nitrogen,  whose  boiling  point 
is  lower  than  that  of  oxygen,  boils  away  first  and  a  residue  of 
nearly  pure  oxygen  is  finally  left.  This  method  is  now  used 
on  a  large  scale  as  the  cheapest  method  of  preparing  oxygen 
for  medicinal  and  commercial  uses. 

Properties  of  Oxygen. — Oxygen  is  a  colorless  and  odorless 
gas.  It  is  slightly  heavier  than  air  and  one-seventh  heavier 
than  nitrogen,  the  gas  which  forms  about  four-fifths  of 
the  volume  of  air. 

The  most  striking  property  of  oxygen  is  its  effect  on 
burning  substances.  A  splinter  of  wood  with  a  live  coal 
at  the  end  will  burst  into  flame  in  the  gas. 

A  piece  of  charcoal  which  has  been  ignited  will  glow 
brilliantly,  and  will  be  surrounded  with  a  small  blue  flame, 
scarcely  visible  at  a  short  distance.  As  the  charcoal  dis- 
appears, an  invisible  gas,  carbon  dioxide,  takes  the  place 
of  the  oxygen.  The  presence  of  the  gas  can  be  shown  by 


BURNING  IN  OXYGEN 


11 


pouring  some  clear  lime  water  into  the -bottle  in  which  the 
charcoal  was  burned.  The  lime  water  will  become  turbid. 
The  precipitate  formed  is  calcium  carbonate,  the  same 
substance  which  is  found  in  marble  and  limestone.  Lime 
is  manufactured  by  heating  limestone  in  a  kiln  till  the 
carbon  dioxide  which  it  contains  is  expelled.  The  lime  is 
"slaked"  by  mixing  it  with  water,  and  lime  water  is  a  solu- 
tion of  this  slaked  lime  in  water.  The  carbon  dioxide  of 
the  candle  finally  converts  the  lime  back  to  the  same  com- 
pound from  which  the  lime  was  made  in  the  lime  kiln. 

Sulfur  burns  in  oxygen  with  a  beautiful  blue  flame  and  the 
oxygen  is  replaced  by  a  colorless  gas  having  a  suffocating 
odor.  This  gas  is  sulfur  dioxide.  It  dis- 
solves in  water,  imparting  to  it  an  acid 
taste,  and  the  solution  will  change  the  color 
of  a  solution  of  litmus  from  blue  to  red. 
Other  acids,  such  as  the  acid  in  vinegar,  in 
a  lemon  or  in  cream  of  tartar,  will  change 
the  color  of  litmus  in  the  same  way.  All 
of  these  acids  have  a  sour  taste. 

Phosphorus  burns  in  oxygen  with  a  very 
brilliant  white  light.  If  the  oxygen  is  dry, 
a  white  powder  settles  on  the  bottom  and 
sides  of  the  bottle  in  which  the  phosphorus 
is  burned.  This  powder  is  phosphorus 
pentoxide.  It  has  a  very  strong  affinity 
for  water  and  will  absorb  water  from  ordinary  air,  com- 
bining with  the  water  and  finally  dissolving  in  it,  giving  the 
solution  an  acid  taste.  The  solution  reddens  blue  litmus. 

An  iron  wire  or  a  coiled  steel  watch  spring,  which  is 
ignited  by  means  of  a  little  sulfur  or  by  a  string  which  has 
been  dipped  in  paraffin  or  sulfur,  will  burn  in  oxygen 
(Fig.  4),  throwing  off  sparks  and  forming  white-hot,  molten 
globules  of  the  magnetic  oxide  of  iron,  which  drop  from 
the  end.  This  oxide  does  not  dissolve  in  water. 


m 


FIG.    4. 


12  COMBUSTION 

Weight  of  Products  Formed.     Indestructibility  of  Matter. 

- — It  is  quite  easy  to  show  that  the  carbon  dioxide,  the  sulfur 
dioxide,  the  phosphorus  pentoxide  and  the  oxide  of  iron 
each  weighs  more  than  the  charcoal,  sulfur,  phosphorus  or 
iron  from  which  it  is  formed.  By  weighing  the  charcoal  or 
other  substance  burned,  the  oxygen  which  is  used  up  in  the 
process  of  burning  and  the  carbon  dioxide  or  other  com- 
pound which  is  formed,  it  has  been  shown  that  the  weight  of 
the  product  formed  by  the  combustion  is  always  exactly 
equal  to  the  weight  of  the  substance  burned  plus  the  weight 
of  the  oxygen  with  which  it  combines.  An  examination 
of  very  many  different  kinds  of  chemical  action  has  shown 
that,  no  matter  how  much  the  substances  which  interact 
with  each  other  may  be  altered  in  their  appearance  and 
properties,  it  has  never  been  possible  to  discover  any 
change  in  the  weight  of  the  materials  involved,  provided 
that  we  consider  the  weights  of  all  the  substances  which 
enter  into  the  action  and  all  of  the  products  formed.  This 
is  known  as  the  law  of  the  indestructibility  of  matter,  and  it  is 
one  of  the  most  important  and  fundamental  of  the  physical 
laws.  It  is  a  law  which  is  always  assumed  in  the  quantita- 
tive study  of  any  chemical  problem. 

Occurrence  of  Oxygen. — Oxygen  is  the  most  abundant  of 
the  elements  and  forms  about  one-half  of  that  part  of  the 
earth  which  we  can  examine.  As  shown  by  Lavoisier's 
experiment,  it  is  found  free  or  uncombined  in  the  air, 
forming  about  one-fifth  of  its  volume.  It  forms  eight- 
ninths  of  the  weight  of  water  and  nearly  one-half  of  the 
weight  of  limestone  and  more  than  one-half  of  the  weight  of 
sand.  It  enters  into  the  composition  of  a  large  part  of  the 
compounds  which  are  known. 

Kindling  Temperature. — Different  substances  vary 
greatly  in  regard  to  the  temperature  at  which  they  will 
take  fire  and  begin  to  burn  rapidly.  Phosphorus  takes  fire 
.at  a  very  low  temperature  and  this  property  is  utilized  in 


SLOW  OXIDATION  13 

matches,  which  may  be  ignited  by  the  heat  produced 
by  gentle  friction.  The  phosphorus  is  used  to  ignite,  in 
turn,  sulfur  or  some  other  easily  combustible  substance 
and  the  latter  in  burning  generates  enough  heat  to  raise 
the  wood  to  its  kindling  temperature.  Somewhat  similar 
phenomena  are  familiar  to  everyone  in  starting  a  fire. 

Heat  of  Combustion. — Various  forms  of  apparatus  have 
been  devised  in  which  it  is  possible  to  burn  charcoal,  sulfur, 
coal,  meaty  bread  and  other  substances  under  conditions 
such  that  the  heat  produced  by  their  combustion  can  be 
absorbed  by  water  and  the  quantity  of  the  heat  generated 
can  be  accurately  measured.  When  this  is  done  it  is  found 
that  the  heat  produced  by  burning  a  kilogram  of  charcoal 
is  nearly  four  times  as  great  as  that  produced  by  burning  a 
kilogram  of  sulfur,  and  about  one-third  greater  than  that 
given  out  in  burning  a  kilogram  of  phosphorus.  A  low 
kindling  temperature  seems  to  indicate  a  strong  affinity 
for  oxygen,  but  a  substance  with  a  high  kindling  temperature 
may  give  more  heat  in  burning. 

Slow  Oxidation. — Meat,  bread,  and  other  articles  of 
food  may  be  burned  in  the  apparatus  which  has  been  re- 
ferred to,  and  the  heat  of  combustion  determined.  In 
this  rapid  burning  the  food  is  converted  chiefly  into  carbon 
dioxide  and  water.  If  these  same  articles  of  food  are  eaten 
the  processes  which  go  on  in  our  bodies  also  convert  them 
into  carbon  dioxide  and  water,  the  oxygen  for  the  purpose 
being  taken  in  as  we  breathe.  As  the  temperature  of  our 
bodies  is  usually  higher  than  that  of  the  surrounding  air,  heat 
must  constantly  escape  from  the  surface  of  our  bodies. 
This  suggests  that  the  heat  necessary  to  maintain  body 
temperature  is  derived  from  the  slow  oxidation  of  the  food 
which  we  eat.  By  means  of  a  complicated  apparatus 
called  a  respiration  calorimeter  the  heat  given  off  from  a 
man's  body  during  several  days  has  been  carefully  measured 
and  compared  with  the  he.at  produced  by  burning  food  of 


14  COMBUSTION 

the  same  kind  as  that  which  he  had  eaten.  In  this  way  it 
has  been  demonstrated  that  food  which  is  eaten  and  then 
oxidized  slowly  in  the  body  gives  just  the  same  quantity  of 
heat  as  the  same  food  when  burned  rapidly. 

Vegetable  and  animal  substances  which  are  exposed  to  the 
action  of  bacteria  in  the  soil  or  elsewhere  are  rapidly  de- 
stroyed by  processes  of  oxidation  somewhat  similar  to  those 
which  go  on  in  our  bodies. 

When  iron  or  steel  is  exposed  to  the  action  of  water  and 
air,  it  becomes  covered  with  rust  by  a  process  of  slow  oxida- 
tion. The  rust  contains  hydrogen  as  well  as  oxygen  and  is 
quite  different  in  composition  from  the  magnetic  oxide  of 
iron  formed  by  the  combustion  of  iron  wire  in  oxygen  gas. 
The  oxide  of  iron  formed  on  the  surface  of  red  hot  iron  ex- 
posed to  the  air  is,  'however,  the  same  as  that  formed  by 
burning  iron. 

Spontaneous  Combustion. — A  mass  of  oily  cotton  waste, 
such  as  is  used  in  cleaning  machinery,  or  a  pile  of  moist 
coal,  will  often  oxidize  slowly  in  the  air  at  ordinary  tempera- 
tures. The  heat  generated  in  this  manner  may  cause  the 
temperature  to  rise  till  the  kindling  temperature  of  the 
material  is  reached,  when  the  mass  will  burst  into  flame. 
This  process  is  known  as  spontaneous  combustion.  The 
spontaneous  combustion  of  coal  may  be  avoided  either  by 
keeping  it  dry  or  by  completely  covering  it  with  water. 

Conservation  of  Energy. — It  has  been  pointed  out  that 
the  oxidation  of  the  food  which  we  eat  furnishes  the  heat 
to  maintain  the  temperature  of  our  bodies.  It  also  fur- 
nishes the  muscular  energy  with  which  we  move.  The 
chemical  energy  of  the  oxygen  and  of  the  elements  or  sub- 
stances which  combine  with  it  is  transformed  into  heat 
energy  and  muscular  energy.  As  is  well  known,  a  part  of 
the  chemical  energy  of  oxygen  and  of  coal,  oil  or  gas  may 
be  converted  into  mechanical  energy  by  a  steam  engine, 
and  the  mechanical  energy  of  the  engine  may  be  converted 


NOMENCLATURE.     OZONE 


15 


to  electrical  energy  by  a  dynamo  or  back  to  heat  by  means 
of  friction.  The  relations  between  these  different  forms 
of  energy  have  been  very  carefully  measured  and  it  has 
been  demonstrated  that  whenever  one  kind  of  energy 
disappears  an  exactly  equivalent  amount  of  some  other 
form  of  energy  takes  its  place.  Energy  cannot  be  created 
or  destroyed  by  any  process  known  to  us.  This  is  the  law 
of  the  conservation  of  energy.  It  might  be  called  the  law 
of  the  indestructibility  of  energy. 

Names  of  Binary  Compounds. — It  will  have  been  noticed 
that  the  compounds  formed  by  the  union  of  oxygen  with 
other  elements  are  called  oxides.  Other  compounds  of  two 
elements  are  given  names  with  the  same  ending, — ide, 
added  to  the  name  of  the  non-metallic  element. .  Com- 
pounds of  sulfur  are  called  sulfides,  compounds  of  chlorine, 
chlorides,  compounds  of  bromine,  bromides.  The  prefixes 
in  such  names  as  dioxide  and  pentoxide  refer  to  the  amount 
of  oxygen  in  the  compound  and  will  be  explained  later. 


FIG.  5. 


Ozone. — If  the  oxygen  is  subjected  to  t^e  action  of  a 
silent  electric  discharge  in  the  apparatus  shown  in  Fig.  5, 
it  acquires  a  peculiar  odor  and  becomes  very  much  more 
active,  combining  readily  with  silver  and  with  other  sub- 
stances which  are  not  affected  by  ordinary  oxygen.  If 
this  changed  oxygen  is  liquefied  by  cooling  it  to  a  very 
low  temperature,  it  is  found  to  be  of  a  dark  blue  color,  and 


16  COMBUSTION 

on  allowing  the  liquefied  gas  to  boil  away  the  last  portions 
of  the  liquid  will  boil  at  a  higher  temperature  than  ordinary 
oxygen  and  will  give  a  gas  which  is  one-half,  heavier  than 
oxygen.  If  this  heavy  gas  is  heated,  however,  it  is  changed 
back  completely  to  oxygen  and  no  other  substance  can  be 
found  in  the  product. 

This  active,  heavy  form  of  oxygen  is  called  "  ozone. " 
It  is  a  powerful  germicide  and  is  sometimes  used  to  sterilize 
water  which  has  been  contaminated  with  disease  germs. 
It  is  possible  that  the  ozone  formed  in  thunder  storms 
has  some  effect  in  purifying  the  air,  but  it  is  rather  doubtful 
if  this  is  of  much  practical  importance. 

Allotropic  Forms  of  Elements. — In  defining  elements  and 
compounds  it  was  implied  that  a  compound  may  always 
be  separated  into  two  or  more  substances  which  we  call 
elements  and  that  these  elements  cannot  be  mutually  converted, 
the  one  into  the  other.  Oxygen  may  be  changed  into  a  mix- 
ture of  ozone  and  oxygen,  but  we  still  call  it  an  element 
because  the  ozone  may  be  completely  changed  back  into 
oxygen  and  there  is  no  change  of  weight  accompanying 
the  process.  Several  other  elements  exist  in  two  or  more 
different  forms.  These  are  called  allotropic  forms.  Thus 
ozone  is  an  allotropic  form  of  oxygen.  Ordinary  phos- 
phorus and  red  phosphorus  are  allotropic  forms  of  that 
element. 

SUMMARY 

Substances  are  changed  in  form  but  not  destroyed  by 
burning. 

The  products  formed  by  burning  wood,  coal  or  other 
substances  weigh  more  than  the  substances  burned. 

Lavoisier  showed  that  mercury,  when  heated  in  the  air, 
takes  up  oxygen  and  that  oxygen  forms  about  one-fifth  of 
the  volume  of  the  air. 

Oxygen  may  be  prepared  by  heating  oxide  of  mercury, 


SUMMARY.     OXYGEN  I/ 

or  by  heating  a  mixture  of  potassium  chlorate  and  manga- 
nese dioxide ;  also  by  the  fractional  distillation  of  liquid  air. 

Carbon,  sulfur,  phosphorus  and  iron  burn  in  oxygen, 
giving  carbon  dioxide,  sulfur  dioxide,  phosphorus  pentoxide 
and  magnetic  oxide  of  iron. 

Carbon  dioxide  gives  a  precipitate  of  calcium  carbonate 
with  lime  water. 

Sulfur  dioxide  and  phosphorus  pentoxide  give  acids  with 
water. 

The  weight  of  the  substance  burned  added  to  the  weight 
of  the  oxygen  with  which  it  combines  is  exactly  equal  to 
the  weight  of  the  compound  formed.  Matter  is  neither 
destroyed  nor  created  by  any  chemical  process. 

Compounds  can  be  formed  from  elements  or  separated 
into  elements,  and  always  weigh  more  than  any  single  ele- 
ment which  they  contain. 

Oxygen  is  the  most  abundant  element  known  in  com- 
pounds on  the  surface  of  the  earth. 

Kindling  temperature  is  the  temperature  at  which  a 
substance  takes  fire  and  burns  rapidly. 

Heat  of  combustion  is  the  heat  generated  by  burning  a 
substance.  It  is  the  same  whether  the  substance  burns 
rapidly  or  oxidizes  slowly. 

Slow  oxidation  may  generate  enough  heat  to  cause  spon- 
taneous combustion. 

Energy  can  neither  be  created  nor  destroyed. 

Elements  may  sometimes  be  changed  to  allotropic  forms, 
but  this  always  occurs  without  change  in  weight. 

Ozone  is  an  active,  allotropic  form  of  oxygen. 

EXERCISES 

1.  Would  the  kindling  temperature  be  the  same  in  oxygen  as 
in  air? 

2.  About  39  per  cent  of  the  weight  of  potassium  chlorate  is 
oxygen.     One  liter  of  oxygen  weighs  1.429  grams.     How  many 

2 


18  COMBUSTION 

grams  of  potassium  chlorate  will  be  required  to  give  one  liter  of 
oxygen? 

3.  Thirty-two  parts  of  oxygen  are  required  to  burn  12  parts  of 
carbon.    How  many  grams  will  be  required  to  burn  one  pound 
(453  grams)   of  pure  charcoal? 

4.  Eighty  grams  of  oxygen  are  required  to  burn  62  grams  of 
phosphorus.     How  many  grams  will  be  required  to  burn  one 
pound? 

6.  Assuming  that  one-fifth  of  the  volume  of  the  air  is  oxygen, 
how  many  liters  of  air  will  be  required  to  burn  a  pound  of  charcoal? 
How  many  liters  will  be  required  to  burn  a  pound  of  phosphorus? 

6.  One  part  of  sulfur  requires  one  part  of  oxygen  for  its  com- 
bustion and  168  grams  of  iron  require  64  grams  of  oxygen.     How 
manj''  grams  of  oxygen  and  how  many  liters  of  air  are  required  to 
burn  a  pound  of  each? 

7.  What  is  the  "oxone"  method  of  preparing  oxygen? 

8.  How  may  oxygen  be  prepared  by  means  of  barium  peroxide 
(Erin's  Process)? 


CHAPTER  III 
HYDROGEN 

Decomposition  of  Water  by  Iron. — If  iron  is  placed  in  an 
iron  or  glass  tube  as  shown  in  the  figure,  and  steam  is  passed 
over  it  while  it  is  heated  red  hot,  the  iron  will  increase  in 
weight  and  will  be  gradually  converted  into  magnetic 
oxide  of  iron  having  the  same  composition  as  the  oxide  of 
iron  formed  when  an  iron  wire  is  burned  in  oxygen.  At  the 


FIG.  6. 

same  time  a  very  light  gas,  hydrogen,  will  collect  in  the 
cylinder  filled  with  water,  which  is  placed  over  the  tube 
through  which  the  excess  of  steam  escapes.  The  experi- 
ment demonstrates  that  water  contains  the  elements  oxygen 
and  hydrogen.  It  can  be  easily  shown  that  water  is  formed 
by  burning  hydrogen  in  oxygen  and  this  completes  the 
proof  that  water  is  a  compound  of  oxygen  and  hydrogen 
and  of  nothing  else. 

Decomposition  of  Water  by  Sodium  or  Potassium.— 
A  piece  of  sodium  wrapped  in  paper  and  thrust  quickly 

19 


20  HYDROGEN 

under  the  mouth  of  an  inverted  cylinder  filled  with  water 
will  act  rapidly  on  the  water,  liberating  hydrogen.  If  the 
sodium  is  thrown  on  the  surface  of  the  water,  the  same  will 
occur,  but  the  hydrogen  will  escape.  If  the  sodium  is  put 
on  a  piece  of  paper  floating  on  the  surface  of  the  water,  to 
prevent  its  rolling  around,  enough  heat  will  be  generated 
so  that  the  liberated  hydrogen  will  take  fire  and  burn. 
The  product  formed  by  the  action  of  the  sodium  on  the 
water  dissolves  in  the  water.  When  the  solution  is  evapo- 
rated a  white  solid  remains,  which  is  composed  of  sodium, 
hydrogen  and  oxygen.  This  is  called, 
because  of  its  composition,  sodium 
hydroxide.  The  name  is  made  up  from 
the  names  of  the  three  elements  which 
it  contains.  The  presence  of  the  sodium 
hydroxide  can  also  be  shown  without 
evaporation,  by  the  action  on  red  litmus 
paper,  which  is  turned  blue  by  the 
solution. 

Potassium  acts  on  water,  in  a  similar 
manner  but  so  much  heat  is  generated 
that  the  hydrogen  catches  fire  without 
the  use  of  the  paper.  The  yellow  color  of  the  flame  of  the 
hydrogen  from  the  sodium  and  the  violet  color  of  that  from 
the,  potassium  are  due  to  the  metals  and  not  to  the  hydrogen. 
Pure  hydrogen  burns  with  an  almost  invisible  flame. 

We  are  familiar  with  many  articles  of  food  which  have  a 
sharp,  acid  taste.  The  most  familiar  of  these  are  vinegar 
which  contains  acetic  acid,  and  lemons  which  contain  citric 
acid.  It  has  been  pointed  out  (p.  11)  that  the  solutions 
formed  by  dissolving  sulfur  dioxide  and  phosphorus  pentox- 
ide  in  water  also  have  an  acid  taste.  In  both  cases  the 
taste  is  not  due  to  the  oxide,  but  to  a  compound  of  the  oxide 
with  water.  Many  other  substances  give  solutions  with 
this  same  characteristic,  acid  taste  and  it  has  been  found 


FIG.  7. 


PREPARATION  OF  HYDROGEN 


21 


that  all  such  substances  contain  hydrogen.  Two  very  common 
and  important  compounds  of  this  sort  are  hydrochloric  acid, 
a  compound  of  hydrogen  and  chlorine,  and  sulfuric  acid, 
a  compound  of  hydrogen  with  sulfur  and  oxygen. 

Hydrogen  from  Acids  and  Metals. — If  a  dilute  solution 
of  these  acids  is  brought  into  contact  with  zinc  or  iron 
the  hydrogen  of  the  acid  is  displaced  by  the  metal  and 
liberated,  very  much  as  the  hydrogen  of 
water  is  displaced  by  sodium  or  potassium. 
If  hydrochloric  acid  and  zinc  are  used,  the 
zinc  combines  with  the  chlorine  to  form 
zinc  chloride.  If  sulfuric  acid  is  used,  the 
zinc  combines  with  the  sulfur  and  oxygen, 
giving  zinc  sulfate.  When  iron  is  used,  the 
compounds  formed  are  ferrous  chloride  and 
ferrous  sulfate. 

Substances  such  as  zinc  chloride  and 
zinc  sulfarte,  ferrous  chloride  and  ferrous 
sulfate,  formed  by  replacing  the  hydrogen 
of  an  acid  by  a  metal,  are  called  salts. 

The  preparation  of  hydrogen  from  zinc  and  dilute  hydro- 
chloric or  sulfuric  acid  may  be  carried  out  by  means  of  the 
simple  apparatus  shown  in  Fig.  8.  Larger  quantities  of 
hydrogen  may  be  obtained  conveniently  by  means  of  a 
Kipp  generator  (Fig.  9).  The  zinc  is  placed  in  the  middle 
bulb  and  the  dilute  acid  is  poured  in  through  the  upper 
bulb,  which  communicates  with  the  lower  one  through  the 
tube  A.  When  the  stopcock  B  is  opened,  the  acid  rises  and 
comes  in  contact  with  the  zinc  in  the  middle  bulb  and  the 
generation  of  hydrogen  begins.  Whenever  the  stopcock 
is  closed  the  pressure  of  the  hydrogen  generated  forces  the 
acid  away  from  the  zinc  and  the  action  ceases  as  soon  as  the 
acid  moistening  the  surface  of  the  zinc  is  exhausted. 

In  using  any  hydrogen  generator,  because  of  the  explosive 
character  of  mixtures  of  hydrogen  and  air,  it  is  necessary 


FIG.  8. 


22 


HYDROGEN 


to  be  careful  not  to  bring  a  flame  near  the  exit  tube  before 
the  air  in  the  generator  is  displaced  by  hydrogen.  Whether 
the  air  has  been  expelled  sufficiently  for  safety  can  be 
determined  by  filling  a  test-tube  with  the  gas  in  the  manner 
shown  in  Fig.  10.  After  filling  the  tube  and  removing  it  to 
a  safe  distance  from  the  exit  tube,  with  the  mouth  of  the 
test-tube  always  held  downward,  a  lighted  match  or  flame 
may  be  applied  to  its  mouth.  If  there  is  a  rather  sharp 


FIG.  9. 


FIG.    10. 


explosion,  extending  up  into  the  tube,  there  is  still  enough 
air  mixed  with  the  hydrogen  to  render  it  dangerous.  If 
there  is  only  a  slight  explosion  and  the  hydrogen  continues 
to  burn  in  the  tube  at  the  surface  where  the  gas  is  in  contact 
with  the  air,  the  hydrogen  is  pure  enough  to  be  safe. 

Properties  of  Hydrogen. — Hydrogen  is  the  lightest  gas 
known.  One  liter  at  the  freezing  point  of  water  (0°), 
and  under  normal  atmospheric  pressure,  equal  to  that  of 
760  mm.  of  mercury,  weighs  0.09  gram,  while  a  liter  of  air 
weighs  1.293  grams,  about  14^  times  .as  much.  A  liter 


REDUCTION  23 

of  oxygen  weighs  1.429  grams,  nearly  16  times  as  much 
as  a  liter  of  hydrogen.  Because  of  its  lightness  hydrogen 
is  used  to  fill  balloons.  Soap-bubbles  filled  with  the  gas 
rise  through  the  air. 

Hydrogen  may  be  liquefied  by  a  process  similar  to  that 
which  has  been  described  for  liquefying  air.  The  com- 
pressed gas  must,  however,  be  cooled  with  liquid  air  before 
it  is  allowed  to  expand  through  the  spiral  tube.  Liquid 
hydrogen  is  very  light  indeed  and  boils  at  a  temperature 
about  65°  below  the  boiling  point  of  liquid  air.  By  causing 
the  liquid  hydrogen  to  boil  under  diminished  pressure  a  part 
of  it  may  be  frozen  to  a  solid,  at  —252.5°. 

The  most  striking  chemical  property  of  hydrogen  is  its 
strong  affinity  for  oxygen.  A  jet  of  hydrogen  will  burn 
quietly  in  air  or  in  oxygen  and  moisture  will  be  deposited 
on  a  cold  surface  held  over  the  flame.  By  collecting  the 
water  formed  by  burning  dry  hydrogen  in  dry  air  and 
determining  its  freezing  point  and  boiling  point,  it  can  be 
shown  that  it  is  really  pure  water,  and  this  proves  that  water 
is  composed  of  two  elements — hydrogen  and  oxygen. 

Mixtures  of  hydrogen  and  air  explode  when  ignited  and 
mixtures  of  hydrogen  and  oxygen  explode  still  more 
violently. 

Occurrence  of  Hydrogen. — The  occurrence  of  hydrogen 
in  water  and  in  acids  has  been  mentioned.  It  is  also  a  con- 
stituent of  practically  all  organic  compounds,  that  is,  of  the 
compounds  of  carbon  which  are  found  in  vegetable  and 
animal  bodies,  in  coal,  petroleum,  natural  gas  and  in  many 
thousands  of  carbon  compounds  which  are  made  in  factories 
for  use  as  dyes  and  medicines.  There  is  a  very  minute 
quantity  of  free  hydrogen  in  the  air,  and  it  is  believed 
that  the  outermost  parts  of  the  atmosphere  are  nearly 
pure  hydrogen. 

Reduction. — Hydrogen  will  not  only  burn  in  air  but  it  will 
also  take  oxygen  away  from  many  oxides  when  these  are 


24  HYDROGEN 

heated  in  a  current  of  the  gas.  Copper  oxide  when  heated 
gives  up  its  oxygen  very  readily  to  hydrogen  which  is  passed 
over  it,  and  is  changed  to  metallic  copper.  Such  a  removal 
of  oxygen  is  called  reduction  and  the  oxide  of  copper  is 
said  to  be  reduced  to  metallic  copper.  At  the  same  time 
the  hydrogen  is  oxidized  to  water.  It  is  evident  that  the 
oxidation  of  the  hydrogen  and  the  reduction  of  the  copper 
oxide  are  opposite  processes.  In  one  oxygen  is  added; 
in  the  other  it  is  removed.  In  almost  all  cases  when  one 
substance  is  reduced  some  reducing  agent  is  at  the  same 
time  oxidized. 

Reversible  Reactions. — If  hydrogen  is  passed  over  heated 
magnetic  oxide  of  iron  a  part  of  the  hydrogen  will  be 
oxidized  to  water  and  the  oxide  of  iron  will  be  slowly  re- 
duced to  metallic  iron.  This  seems  rather  surprising 
because  it  has  been  stated  before  that  when  steam  is 
passed  over  heated  iron  the  iron  is  oxidized  and  hydrogen  is 
liberated.  It  is  natural,  at  first,  to  think  that  the  oxidation 
of  the  iron  by  the  steam  takes  place  because  the  affinity  of 
the  iron  for  oxygen  is  greater  than  that  of  hydrogen. 
When  we  discover,  however,  that  hydrogen  can  take  oxygen 
away  from  oxide  of  iron  it  becomes  evident  that  this 
simple  view  of  chemical  affinity  cannot  be  true.  Since 
the  iron  is  oxidized  when  the  steam  is  in  excess  and  the 
hydrogen  is  constantly  removed  from  the  point  of  action, 
while  the  oxide  is  reduced  when  the  hydrogen  is  in  excess 
and  the  steam  is  all  of  the  time  carried  away  by  the  current 
of  the  gas,  it  must  be  that  the  reaction  is  reversible.  As 
long  as  steam,  hydrogen,  iron  and  oxide  of  iron  are  all  present 
the  action  is  all  of  the  time  going  in  both  directions,  but 
if  the  steam  is  in  excess  the  action  goes  faster  in  the  direction 
toward  the  formation  of  oxide  of  iron  and  hydrogen  and 
it  may  finally  become  complete  in  that  direction.  If  the 
hydrogen  is  in  excess,  it  goes  faster  in  the  other  direction 
and  the  oxide  may  be  completely  reduced  to  metallic  iron. 


OXYHYDROGEN  BLOWPIPE  25 

Oxyhydrogen  Blowpipe. — This  is  an  instrument  by 
means  of  which  oxygen  and  hydrogen  may  be  brought 
together  and  burned,  giving  an  intensely  hot  flame.  The 
temperature  may  be  as  high  as  2000°  to  25000.1  The  tem- 
perature of  the  flame  is  limited  by  the  fact  that  the 
combination  of  the  oxygen  and  hydrogen  is  a  reversible 
reaction.  At  moderate  temperatures  and  even  at  a  tem- 
perature of  2000°  the  reaction  goes  almost  exclusively  in 
the  direction  toward  the  formation  of  water.  At  very  high 


FIG.   11. 

temperatures  water  is  partly  decomposed  into  oxygen  and 
hydrogen  and  at  such  temperatures  there  will  be  a  mixture 
of  oxygen,  hydrogen  and  water.  As  the  temperature 
rises  the  proportion  of  water  in  the  mixture  will  decrease, 
and,  since  the  heat  comes  from  the  union  of  the  oxygen  and 
hydrogen,  the  temperature  which  can  be  obtained  by  burn- 
ing the  mixed  gases  is  limited  because  at  very  high  tem- 
peratures a  considerable  part  of  the  oxygen  and  hydrogen 
remain  uncombined.  At  3700°  a  mixture  of  oxygen, 
hydrogen  and  water  will  contain  only  60  per  cent  of  its 
weight  as  water,  40  per  cent  of  the  oxygen  and  hydrogen 
remaining  uncombined.  The  temperature  of  the  oxy- 
hydrogen  flame  is  much  lower  than  this.  The  temperature 
of  an  electric  arc  between  carbon  poles  is  estimated  at  about 

1  All  temperatures  are  given  in  Centigrade  degrees. 


26  HYDROGEN 

3600°,  while  the  temperature  of  an  open  hearth  steel  furnace 
is  only  1500°  to  1700°. 

Platinum,  which  has  a  melting  point  of  1755°,  melts 
readily  in  the  oxyhydrogen  flame.  Iron  takes  fire  and  burns, 
throwing  off  brilliant  sparks. 

The  flame  alone  gives  very  little  light,  but  if  a  piece  of 
lime  is  placed  in  the  flame  it  glows  intensely  and  gives  a 
very  brilliant  light,  called  variously  the  "lime  light," 
"calcium  light''  or  "Drummond  light."  This  is  often 
used  for  illumination  in  stereopticons  but  has  been  mostly 
replaced  by  electric  lights.  The  oxyhydrogen  blowpipe 
is  used  to  melt  together  the  edges  of  the  leaden  plates  used 
in  making  the  chambers  for  the  manufacture  of  sulfuric 
acid.  A  similar  blowpipe  in  which  acetylene  is  used  in 
place  of  hydrogen  and  which  gives  a  still  hotter  flame  is  now 
extensively  used  for  cutting  steel  or  iron  bars  or  plates. 
The  flame  will  melt  its  way  through  the  metal  very  rapidly. 
The  blowpipe  is  also  used  in  blacksmith  shops  and  garages 
for  welding. 

SUMMARY 

Water  may  be  decomposed  by  red-hot  iron,  giving  mag- 
netic oxide  of  iron  and  hydrogen. 

Sodium  and  water  give  hydrogen  and  sodium  hydroxide. 
Potassium  gives  hydrogen  and  potassium  hydroxide. 

Acids,  especially  hydrochloric  acid  and  sulfuric  acid, 
contain  hydrogen  which  can  be  displaced  by  zinc  or  iron, 
giving  the  salts,  zinc  chloride  or  sulfate  or  ferrous  chloride 
or  sulfate. 

Hydrogen  is  the  lightest  gas  known.  It  may  be  lique- 
fied and  frozen  at  very  low  temperatures.  It  burns  in  air, 
forming  water. 

Hydrogen  occurs  in  water,  acids,  organic  compounds  and, 
in  minute  quantities,  in  air. 

Hydrogen  reduces  hot  copper  oxide  to  metallic  copper. 


SUMMARY.     HYDROGEN  27 

It  also  reduces  magnetic  oxide  of  iron  but  the  reaction  is 
reversible. 

The  union  of  oxygen  and  hydrogen  is  also  reversible  at 
high  temperatures  and  this  limits  the  temperature  of  the 
oxyhydrogen  blowpipe. 

The  oxyacetylene  flame  is  used  for  cutting  steel  and  iron. 

EXERCISES 

1.  One  part  by  weight  of  hydrogen  requires  eight  parts  by  weight 
of  oxygen  for  its  combustion.     How  many  liters  of  oxygen  will  be 
required  to  burn  a  liter  of  hydrogen?     How  many  liters  of  air? 
Suggestion:  Find  first  from  the  text  the  weight  of  a  liter  of  hydro- 
gen and  of  a  liter  of  oxygen  and  the  proportion  of  oxygen  in  the 
air.     Answers  are  to  be  given  in  round  numbers  or  with  two  or 
three  significant  figures,   not  with  long  decimals. 

2.  Sugar  contains  44  per  cent  of  carbon  and  6.7  per  cent  of 
hydrogen.     How  many  grams  of  oxygen  will  be  required  to  burn  a 
pound  of  sugar?     How  many  liters  of  air? 

3.  How  many  grams  of  water  and  how  many  grams  of  carbon 
dioxide  will  be  formed  by  burning  a  pound  of  sugar? 

4.  From  a  mixture  of  liquid  air  and  liquid  hydrogen  in  what 
order  would  the  elements  distil  away?     What  would  be  the  effect 
of  applying  a  flame  to  such  a  mixture? 

5.  What  is  the  "Hydione"  method  of  preparing  hydrogen? 

6.  How  is  the  hydrogen  to  fill  Zeppelins  prepared? 


CHAPTER  IV 

WEIGHTS    AND    MEASURES,    MOLECULAR    THEORY, 
LAWS  OF  GASES 

Weights  and  Measures  in  Scientific  Use. — For  scientific 
purposes  the  weights  and  measures  of  the  metric  system  are 
used,  almost  exclusively. 

The  measure  of  length  is  the  meter,  which  is  39.3709 
inches.  Its  subdivisions  are  decimeters,  centimeters,  and 
millimeters,  which  are,  respectively,  tenths,  hundredths 
and  thousandths  of  a  meter. 

The  measure  of  weight  is  the  gram,  which  is  almost,  but 
not  exactly,  the  weight  of  one  cubic  centimeter  of  water 
taken  at  its  maximum  density  (i.e.,  at  4°).  Its  most  com- 
mon subdivision  is  the  milligram,  one-thousandth  of  a 
gram,  and  the  most  common  multiple  is  the  kilogram,  which 
is  one  thousand  grams. 

The  measure  of  volume  is  the  liter ;  which  is  the  volume 
occupied  by  one  kilogram  of  water  at  its  maximum  density. 
It  is  almost,  but  not  exactly,  one  cubic  decimeter.  The 
most  common  division  is  the  cubic  centimeter,  one-thou- 
sandth of  a  liter. 

The  following  approximate  equivalents  are  sometimes 
convenient : 

One  meter  is  a  little  more  than  a  yard  (39.3709  inches). 

One  millimeter  is  almost  ^5  of  an  inch  (0.03937  inch). 

One  gram  is  about  15  grains  (15.432  grains). 

One  kilogram  is  about  2J£  pounds  (2.2046  pounds). 

One  liter  is  a  little  less  than  a  quart  (0.8836  quart). 

One  cubic  centimeter  is  about  one-thirtieth  of  a  fluid 
ounce. 

28 


MOLECULAR  THEORY  29 

Molecular  Theory. — Very  many  and  very  diverse  phe- 
nomena are  most  easily  explained  by  supposing  that  all 
material  objects  are  composed  of  very  minute  particles, 
which  are  called  molecules.1  In  a  proper  scientific  treat- 
ment the  various  facts  and  relations  which  have  led  to  this 
conclusion  should  be  presented  as  a  foundation  for  this 
theory.  Practically,  however,  these  facts  are  much  more 
easily  understood  and  remembered  if  they  are  presented  in 
connection  with  the  theory,  and  the  theory  is  so  useful  that 
it  seems  best  to  give  an  outline  of  it  here.  No  student 
should,  however,  be  satisfied  to  accept  the  theory  on  the 
authority  of  a  book  or  a  teacher.  On  the  contrary,  the 
student  should  hold  the  theory  in  the  earlier  months  of  his 
study  as  something  which  has  not  been  fully  demonstrated, 
and  he  should  return  to  the  theory  over  and  over  again  as 
new  facts  related  to  it  are  learned. 

Solids  retain  their  shape  unless  they  are  subjected  to 
some  force  great  enough  to  bend  or  break  them.  Liquids 
flow  to  the  bottom  of  the  vessels  which  contain  them  and 
present  at  their  top  a  level  surface,  which  is  sharply  sepa- 
rated from  the  gas  or  vapor  above.  Gases  give  no  similar 
surface  but  completely  fill  the  vessel  in  which  they  are  con- 
tained. It  seems  evident  from  these  properties  that  the 
molecules  of  solids  are  held  rather  rigidly  in  position  by 
attractive  forces  between  them.  The  molecules  of  liquids, 
on  the  other  hand,  must  glide  or  slip  easily  over  each  other 
but  are  still  held  together  by  attractive  forces.  The  first 
and  most  natural  opinion  about  gases  was  that  the  molecules 
repel  each  other  and  separate  for  this  reason.  Another 
explanation  of  this  property  is  given  below. 

Diffusion  of  Gases.— If  the  mouths  of  two  narrow  cylin- 
ders containing  air  and  hydrogen  are  brought  together 

1  According  to  the  molecular  theory  molecules  are  the  smallest  particles 
of  any  substance  which  can  exist  alone.  Atoms  are  the  smallest  particles 
of  an  element.  They  may  be  identical  with  the  molecules  of  the  free 
element  but  usually  are  not  (p.  136). 


30 


WEIGHTS  AND  MEASURES 


(Fig.   12)   with  the  cylinder  containing  hydrogen  above, 

although  the  air  is  14)^  times  as  heavy  as  the  hydrogen, 

some  of  the  air  will  make  its  way  upward  into  the  hydrogen 
and  some  of  the  hydrogen  will  make  its  way 
downward  into  the  air.  This  can  be  shown  by 
testing  the  gas  in  each  cylinder  with  a  flame. 
Each  cylinder  will  be  found  to  contain  an  ex- 
plosive mixture  of  air  and  hydrogen.  Such  a 
process  by  which  two  gases  or  two  liquids  in 
contact  with  each  other  mix  is  called  diffusion. 
Two  liquids  in  contact  will  diffuse  into  each 
other  only  when  they  are  mutually  soluble. 
Two  gases  in  contact  will  always  diffuse,  even 
though  the  density  of  one  of  the  gases  be  one 
hundred  times  that  of  the  other  and  no  matter 
how  different  or  how  in- 
soluble in  each  other  the 
same  substances  may  be 

when  they  are  in  the  liquid  state. 
If  a  cylinder  of  porous  porcelain,  with 

openings  so  fine  that  pressure  will  cause 

a  gas  to  pass  through  them  only  very 

slowly,  is  fitted  with  a  rubber  stopper 

and   connected   with   a  bulb  and  bent 

tube  containing  water  as  shown  in  Fig. 

13,  when  a  beaker  filled  with  hydrogen 

is  placed  over  the  cylinder  the  pressure 

within  will  suddenly  increase  and  force 

water  out  of  the  tube  in  a  jet.     This 

shows   that   hydrogen  passes  through 

the  walls  of  the  cylinder  to  the  interior. 

It  can  be  shown  in  this  case,  also,  that 

some  air  passes  out  through  the  walls 

of  the  cylinder  but  the  light  hydrogen  passes  through  the 

fine  openings  very  much  more  rapidly  than  the  air. 


FIG.   12. 


KINETIC  THEORY  OF  GASES 

Kinetic  Theory  of  Gases. — Gases,  under  ordinary  c 
ditions,  are  much  less  dense  than  solids  or  liquids.  If 
quart  of  water  is  changed  to  steam  at  atmospheric  pressure, 
there  will  be  more  than  1600  quarts  of  steam.  It  seems 
improbable  that  the  particles  (molecules)  of  water  increase 
very  much  in  size  when  the  water  is  changed  to  steam. 
If  they  do  not,  it  must  be  that  the  steam  has  1600,  or  more, 
quarts  of  empty  space  for  every  quart  of  space  actually 
filled  by  molecules  of  water.  This  and  other  facts,  some 
of  which  will  be  given  below,  have  led  to  the  proposal  and 
development  of  the  kinetic  theory  of  gases,  a  theory  which 
explains  many  of  the  properties  of  gases  by  considering  the 
motions  of  the  molecules.  According  to  this  theory  the 
molecules  of  a  gas  are  kept  apart  because  they  are  in  very 
rapid  motion  and  because  they  rebound  like  elastic  balls 
whenever  they  hit  one  another  or  when  they  hit  any  solid 
substance.  The  pressure  exerted  by  a  gas  is  evidently 
caused  by  the  bombardment  of  any  surface  with  which 
it  is  in  contact  by  the  molecules  of  the  gas,  just  as  pressure 
would  be  exerted  on  a  wall  if  a  large  number  of  elastic  balls 
were  constantly  thrown  against  it. 

It  is  evident  that  according  to  this  theory  a  gas  expands 
and  completely  fills  any  space  which  is  given  to  it,  not  be- 
cause of  any  repulsion  between  the  molecules  of  the  gas, 
but  because  the  motion  of  the  molecules  causes  them  to 
fly  out  and  fill  any  space  at  the  side,  if  there  is  no  wall  to 
restrain  them.  Indeed,  many  facts  which  have  been  learned 
about  gases  indicate  that  when  the  molecules  come  close 
together  in  their  collisions  there  is  an  attraction  between 
them  and  that  it  is  largely  for  this  reason  that  most  gases 
do  not  obey  exactly  the  laws  of  temperature  and  pressure 
which  are  given  in  the  latter  part  of  this  chapter. 

The  kinetic  theory  gives  a  very  satisfactory  explanation 
of  the  diffusion  of  gases.  When  two  gases  are  brought  into 
contact  the  moving  molecules  of  one  can  readily  shoot  into 


WEIGHTS  AND  MEASURES 

spaces  between  the  molecules  of  the  other  and  some  of 
molecules  which  shoot  in  will  rebound  after  collision  in 
a  manner  as  to  fly  further  into  the  mass  of  the,  other 
gas.  As  the  velocities  of  the  molecules  are  very  great  the 
gases  will  rapidly  become  mixed  at  their  surfaces  and  grad- 
ually throughout  their  whole  mass.  This  process  of  diffu- 
sion can  be  made  apparent  to  the  eye  by  putting  a  little 
bromine  in  the  bottom  of  a  tall  cylinder  containing  air. 
Bromine  vapor  is  colored  and  its  gradual  diffusion  into  the 
air  can  easily  be  seen. 

Collision   Between   Elastic   Bodies. — When  two  elastic 
balls  of  the  same  size,  going  in  opposite  directions,  meet 


FIG.  14. 

squarely,  each  rebounds  with  a  velocity  such  that  its 
energy  is  the  same  as  that  of  the  other  ball  before  contact. 
Thus  if  two  balls,  A  and  B,  have  the  same  weight,  if  A  is 
raised  a  certain  distance  on  one  side  (Fig.  14)  and  allowed 
to  fall  it  will  stop  almost  completely  on  hitting  5,  while  the 
latter  will  rise  to  almost  the  same  distance  on  the  other  side. 
If  A  is  raised  through  an  arc  of  60°  while  B  is  raised  only  30°, 


NUMBER  OF  MOLECULES  33 

after  impact  B  will  recoil  nearly  60°  while  A  will  recoil 
only  about  30°.  If  one.  ball  is  heavier  than  the  other,  the 
lighter  ball  will  recoil  with  a  greater  velocity  than  that 
of  the  heavy  one  when  struck  by  it  and  the  heavy  ball 
will  have  a  smaller  velocity  than  that  of  the  lighter  one  if 
the  heavy  ball  is  at  rest  and  is  struck  by  the  lighter  one. 

Number  of  Molecules  in  the  Same  Volume  of  Different 
Gases. — Applying  these  principles  to  the  consideration 
of  a  mixed  gas  containing  molecules  of  different  weights, 
it  has  been  shown  that  the  average  velocities  of  the  mole- 
cules of  the  different  constituents  of  such  a  gas  must  vary 
inversely  as  the  square  roots  of  the  weights  of  the  mole- 
cules. Also,  the  greater  velocity  of  the  lighter  molecules 
will  keep  them  just  as  far  apart  and  give  the  same  pressure 
as  the  slower  velocity  of  the  heavy  ones.  It  also  follows 
that  when  the  pressure  is  the  same  there  must  be  the  same  num- 
ber of  molecules  in  a  given  volume  of  a  light  gas  as  in  the  same 
volume  of  a  heavy  one.  The  differences  in  the  velocities  of 
light  and  heavy  molecules  explain  very  satisfactorily  the 
diffusion  of  hydrogen  through  the  porous  wall,  and  that 
experiment,  in  turn,  gives  strong  support  for  the  kinetic 
theory.  Molecules  of  oxygen,  according  to  the  theory, 
must  be  about  16  times  as  heavy  as  molecules  of  hydrogen 
and  will,  therefore,  fly  only  one-fourth  as  fast.  As  there  are 
the  same  number  of  molecules  on  both  sides,  four  times  as 
many  of  the  fast  hydrogen  molecules  will  hit  the  small 
openings  and  fly  through  them  in  a  given  time  as  of  the 
slower  oxygen  molecules. 

Number  of  Molecules  in  One  Cubic  Centimeter  of  a  Gas. 
— Several  different  methods  have  been  found  for  estimating 
the  number  of  molecules  in  a  cubic  centimeter  of  a  gas 
under  standard  conditions.  The  most  accurate  of  these 
indicate  that  the  number  is  about  2.71  X  1019  (or  27,100,- 
000,  000,  000,  000,  000).  The  average  velocity  of  hydro- 
gen molecules  at  ordinary  temperatures  is  somewhat  more 


34 


WEIGHTS  AND  MEASURES 


than  a  mile  a  second.     That  of  oxygen  molecules  is  a  little 
more  than  a  quarter  of  a  mile  a  second. 

The  important  conclusion  that  there  are  the  same  num- 
ber of  molecules  in  equal  volumes  of  different  gases  under 
the  same  conditions  of  temperature  and  pressure  was 
TEMPERA-  VOLUME  reached  from  a  consideration 
of  chemical  facts  many  years 
373  cc.  before  tne  kinetic  theory  of 
gases  was  proposed.  We  shall 
283  cc.  have  occasion  to  come  back  to 


ABSOLUTE 
TEMPERATURE 


373C 


283° 
273° 


100C 


10C 


173C 


-100°  --  173cc. 


73< 


Oc 


-200< 


-273° 

FIG.  15. 


0°  "I-  273  cc.  this  later  (P-  135). 

Temperature. — Tempera- 
tures are  measured,  for  scien- 
tific uses,  with  the  Centigrade 
thermometer.  The  freezing 
point  of  water  is  taken  as  0° 
and  the  boiling  point  of  water, 
under  a  pressure  of  one  atmos- 
phere (760  millimeters  of  mer- 
cury), is  taken  as  100°. 

Absolute  Zero. — If  the 
pressure  is  kept  constant, 
hydrogen  expands,  when  it  is 
heated,  at  the  rate  of  3/273  of 
its  volume  at  zero  for  each 
degree.  If  it  is  cooled,  it  con- 
tracts at  the  same  rate,  ^73 
of  its  volume  at  zero,  for  each 
degree.  Evidently  if  it  con- 
tinued to  contract  at  the  same  rate  it  would  disappear  at 
a  temperature  of  —  273°.  It  does  not  continue  to  contract 
at  the  same  rate,,  of  course,  because  it  condenses  to  a  liquid 
at  —252.5°  and  its  expansion  and  contraction  are  then  at 
a  very  different  rate.  Other  gases,  however,  expand  and 
contract,  when  heated  or  cooled,  at  almost  the  same  rate 


73  cc. 


LAW  OF  CHARLES  35 

as  hydrogen.  This  fact,  that  all  gases  expand  and  con- 
tract at  nearly  the  same  rate  for  changes  of  temperature,  has 
made  it  seem  appropriate  to  call  —273°  absolute  zero,  and 
to  call  temperatures  reckoned  from  that  point  absolute 
temperatures.  Before  this  temperature  is  reached  all  gases 
become  liquid,  though  the  boiling  point  of  helium  is  only 
4.5°  absolute. 

The  relations  between  ordinary  and  absolute  tempera- 
tures and  between  these  and  the  volume  of  a  gas  which  ex- 
pands or  contracts  under  constant  pressure  will  be  seen 
from  Fig.  15. 

There  are'  many  reasons  for  thinking  that  the  absolute 
zero  is  not  merely  a  convenient  fiction  or  theory  based  on 
the  conduct  of  gases.  It  seems  to  be  an  actual  limit  of 
temperature  which  is  a  real  starting  point  for  many  differ- 
ent phenomena  and  which  can  never  be  reached  by  experi- 
ment. The  lowest  temperature  so  far  obtained  is  estimated 
at  about  -270°  or  3°  absolute. 

Law  of  Charles. — The  effect  of  a  change  in  temperature 
upon  a  gas  under  constant  pressure  is  conveniently  stated 
in  the  law  of  Charles:  The  volume  of  a  gas  varies  directly 
as  the  absolute  temperature.  This  may  also  be  stated  as  a 
proportion : 

V  :  V  : :  T  :  T 

where  V  and  V '  represent  two  volumes  of  the  same  gas  at  the 
temperatures  T  and  T'.  Thus  273  cc.  (cubic  centimeters) 
of  a  gas  at  0°  will  become  283  cc.  at  10°  (=  273°  +  10°  or 
283°  absolute)  or  263  cc.  at  -10°  (=  263°  absolute). 

A  very  common  problem  in  dealing  with  gases  is  to  find 
the  volume  which  a  gas  measured  at  some  given  temperature 
would  occupy  if  cooled  or  heated  to  0°.     Such  a  problem 
may  be  easily  solved  by  means  of  the  proportion: 
VQ:V::T<>(=  273°):T(  =  t°  +  273°) 

V.-F 


36 


WEIGHTS  AND  MEASURES 


,r 

1 


Such  a  formula  should  never  be  used  mechanically,  but 
the  student  should  merely  remember  that  the  volume 
measured  is  to  be  multiplied  by  a  fraction  whose  numerator 
and  denominator  are  the  two  absolute  temperatures  con- 
cerned and  that  a  gas  expands  when  heated  and  contracts 
when  cooled. 

Pressures. — The  pres- 
sure of  the  air  is  most  easily 
and  accurately  determined 
by  measuring  by  means  of 
a  mercury  barometer  the 
height  of  the  column  of 
mercury  which  will  balance 
it.  For  this  reason  the 
pressure  of  a  gas  is  usually 
given  in  millimeters  of  mer- 
cury, meaning  the  height,  in 
millimeters,  of  the  column 
of  mercury  which  the 
pressure  of  the  gas  will 
sustain,  or  balance.  On 
1  the  average,  at  sea-level, 

the    pressure    of    the    at-- 
]»    mosphere    will    sustain    a 

Barometer  ^J  ^J  ^J  column  of  mercury  760 
3  C  mm.  high,  and  this  is  taken 
as  the  standard  atmos- 
pheric pressure. 
Law  of  Boyle. — When  the  pressure  exerted  on  a  gas  is 
increased  or  diminished  the  volume  of  the  gas  decreases  or 
increases  in  inverse  proportion  to  the  pressure.  If  we  re- 
call the  kinetic  theory  it  will  be  seen  at  once  that  this 
should  be  true.  If  a  gas  is  compressed  to  one-half  its 
volume,  twice  as  many  molecules  must  strike  a  square 
centimeter  of  surface  in  a  given  time.  This  must  cause 


FIG.  16. 


VOLUME  UNDER  STANDARD  CONDITIONS          37 

twice  the  pressure  on  the  surface,  if  the  pressure  is  due 
to  the  bombardment  of  the  surface  by  the  molecules. 

The  truth  of  this  law  may  be  demonstrated  by  means  of 
the  apparatus  shown  in  Fig.  16.  If  the  mercury  is  at  the 
same  level  at  both  arms  of  the  U-tube  in  A,  it  is  evident 
that  the  air  in  the  graduated  part  will  be  at  atmospheric 
pressure  (760  mm.  at  standard  pressure).  If  mercury 
is  drawn  out  from  the  longer  arm,  as  shown  in  B,  till  the 
level  in  that  arm  is  one-half  the  height  of  the  barometer 
column  (380  mm.)  lower  than  in  the  other  arm,  the  pressure 
in  the  graduated  tube  will  be  decreased  to  one-half  of 
an  atmosphere  and  the  volume  of  the  air  will  be  twice 
as  great  as  before.  On  the  other  hand,  if  mercury  is  poured 
into  the  long  arm  till  the  level  is  as  high  as  the  length  of 
the  barometer  column  (760  mm.)  above  the  level  in  the 
graduated  tube,  as  shown  in  C,  the  pressure  will  be  doubled 
and  the  volume  of  the  air  will  be  decreased  to  one-half. 

The  law  of  Boyle  may  be  conveniently  stated  by  the 
proportion  : 

V:V'::P':PorVP  =  V'P'  or  V  =  V'^ 

For  use  in  finding  the  volume  of  a  gas  under  atmospheric 
pressure  when  the  volume  under  some  other  pressure  is 
known  this  becomes 


Reduction  of  the  Volume  of  a  Gas  to  its  Volume  under 
Standard  Conditions.  —  For  the  comparison  of  the  properties 
of  different  gases  and  for  many  other  purposes  it  is  con- 
venient to  have  standard  conditions  of  temperature  and 
pressure.  The  standard  conditions  which  have  been  chosen 
are  0°  Centigrade  for  temperature  and  760  mm.  for  pressure. 
When  the  volume  of  a  gas  is  measured  at  some  other 
temperature  and  pressure,  as  is  usually  convenient,  the 
volume  under  standard  conditions  is  found  by  multiply- 


38 


WEIGHTS  AND  MEASURES 


ing  the  observed  volume  by  two  fractions  in  which  the 
numerator  and  denominator  of  one  are  the  two  absolute 
temperatures  concerned  and  the  numerator  and  denomi- 
nator of  the  other  are  the  two  pressures.  The  formula  is: 

'  0°  and  760   =:     •    jari  X     ^T~ 

To  illustrate,  suppose  15.2  cc.  of  gas  are  measured  at  20° 
and  under  a  pressure  of  735  mm.     The  volume  at  0°  and 

735       273 
760  mm.  will  be  15.2  X  ^r^  X 


As  remarked  of  the  formula  for  the  law  of  Charles,  such 
a  formula  should  never  be  used  mechanically,  but  always 
with  a  clear  understanding  of  its  meaning  and  remembering 
that  an  increase  in  temperature  innreaafis  tfrfi 


an  increase  in  pressure  decreases  the  volume  of  a  gas. 
^Determination  of  the  Weight  of  a  Liter  of  a  Gas.— 
Although  gases  are  very  light  in  comparison  with  solids 


Toolrfxsrnfj 


FlG.     17. 


and  liquids,  they  are  heavy  enough  to  be  weighed  with  a 
fair  degree  of  accuracy  on  a  sensitive  balance.  For  this 
purpose  the  bulb  shown  in  Fig.  17  is  exhausted  till  the 
mercury  in  the  manometer  connected  with  it  stands  level 
in  the  two  arms.  The  stopcock  is  then  closed  and  the 


WEIGHING  A  GAS  39 

bulb  is  weighed.  It  is  then  filled  with  the  gas  to  be  ex- 
amined, the  temperature  and  barometer  are  read  and  the 
bulb  is  weighed  again.  The  increase  in  weight  is,  of  course, 
the  weight  of  the  gas  contained  in  the  bulb.  The  volume 
of  the  bulb  is  determined  by  weighing  it  empty  and  full 
of  water.  Having  in  this  way  determined  the  weight  of  a 
known  volume  of  the  gas  under  a  known  temperature  and 
pressure,  it  is  easy  to  calculate  the  volume  which  the  same 
gas  would  occupy  at  zero  and  atmospheric  pressure,  and' 
from  this  the  weight  of  one  liter  of  the  gas  under  standard 
Conditions. 

The  process  will  be  more  clearly  understood  by  calculating 
the  weight  of  a  liter  of  air  under  standard  conditions  from 
the  following  data: 

Weight  of  bulb  full  of  water 178  grams 

Weight  of  empty  bulb 64  grams 

Weight  of  water : 114  grams 

Hence  the  volume  of  the  bulb  is  114  cc. 

Weight  of  the  bulb  full  of  air  at  23°  and 

742  mm 64.0000 

Weight  of  evacuated  bulb 63 . 8668 

Weight  of  air  contained  in  the  bulb  at 
23°  and  742  mm 0. 1332 

In  solving  this  problem  the  volume  of  air  in  the  bulb 
is  first  reduced  to  its  volume  under  standard  conditions. 
Remembering  that  a  liter  is  1000  cc.  the  calculation  of  the 
weight  of  a  liter  of  air  is  easy. 

SUMMARY 

The  measures  and  weights  most  often  used  in  scientific 
work  are  the  meter  and  millimeter,  gram,  milligram  and 
kilogram,  and  the  liter  and  cubic  centimeter. 

According  to  the  molecular  theory  substances  are  com- 
posed of  very  small  particles  called  molecules,  which  are 


40  WEIGHTS  AND  MEASURES 

held  together  rather  rigidly  in  solids,  slide  over  each  other 
but  are  still  held  together  by  their  attractions  in  liquids, 
and  are  comparative^  far  apart  in  gases. 
'  Gases  in  contact  always  diffuse  into  each  other. 

The  diffusion  and  pressure  of  gases  are  explained  by  the 
kinetic  theory,  which  supposes  that  the  molecules  of  gases 
are  elastic  bodies,  in  rapid  motion. 

Elastic  bodies  in  collision  exchange  their  energy  of  motion. 

Equal  volumes  of  different  gases  contain  the  same  number 
of  molecules  under  the  same  conditions  of  temperature  and 
pressure. 

One  cubic  centimeter  of  a  gas  under  standard  conditions 
contains  2.71  X  1019  molecules. 

Temperatures  are  fixed  with  reference  to  the  freezing 
point  and  boiling  point  of  water  with  100°  between. 

An  absolute  temperature  is  the  temperature  measured 
from  273°  below  zero. 

The  volume  of  a  gas  varies  directly  as  the  absolute  tem- 
perature (law  of  Charles) . 

Pressures  of  gases  are  measured  by  the  height  in  milli- 
meters of  the  column  of  mercury  the  pressure  will  sustain. 
Atmospheric  pressure  is  taken  as  760  mm. 

The  volume  of  a  gas  varies  inversely  as  the  pressure  (law 
of  Boyle). 

The  two  laws  may  be  used  in  calculating  volumes  of 

P'       T 
gases  by  means  of  the  formula  V  =  V  X  -p  X  TJT, 

A  gas  may  be  weighed  in  a  bulb  which  has  been  evacuated 
and  weighed  empty  and  then  full  of  the  gas. 

EXERCISES 

1.  What  will   be  the  volume  under  standard  conditions  of 
37.5  cc.  of  a  gas  measured  at  22°  and  a  pressure  of  735  mm? 

2.  What  will  be  the  volume  at  25°  and  730  mm.  of  44.2  cc.  of 
gas  measured  at  27°  and  770  mm.? 


EXERCISES.     GASES  41 

3.  If  one  liter  of  air  measured  under  standard  conditions  is 
brought  to  a  temperature  of  12°  and  a  pressure  of  620  mm.,  what 
will  be  its  volume? 

4.  What  will  be  the  weight  of  a  liter  of  air  at  12°  and  620  mm. 
if  it  weighs  1.293  grams  under  standard  conditions?     What  will 
be  the  weight  01  a  liter  of  hydrogen  at  the  same  temperature 
and  pressure,  if  it  weighs  0.09  gram  under  standard  conditions? 

6.  What  will  be  the  lifting  power  of  a  balloon  having  a  volume 
of  100  cubic  meters  and  filled  with  hydrogen,  if  it  is  one  mile1 
above  sea-level  and  the  temperature  is  12°  and  the  pressure 
620  mm.? 

6.  Are  there  any  other  gases  besides  hydrogen  which  might  be 
used  to  fill  balloons? 

7.  Since  a  gas  expands  and  fills  completely  any  space  which  is 
given  to  it,  why  does  not  the  atmosphere  expand  and  fill  the  space 
between  it  and  the  sun? 

1  The  height  is  not  used  in  the  calculation.  It  is  given  as  the  height 
above  sea-level  where  the  pressure  is  approximately  620  mm. 


CHAPTER  V 

COMPOSITION   OF   WATER.    LAWS   OF  COMPOSITION 
BY  WEIGHT.     THE  ATOMIC  THEORY 

Analysis.  Synthesis. — Two  general  methods,  which 
have  been  referred  to  briefly,  are  used  in  determining  the 
composition  of  substances:  analysis  and  synthesis.  In 
analysis  the  substance  is  decomposed  into  its  elements,  or, 
more  often,  the  elements  of  which  it  is  composed  are  con- 
verted into  other  compounds  which  can  be  identified  and 
whose  composition  is  known.  The  decomposition  of  oxide 
of  mercury  into  mercury  and  oxygen  is  the  simplest  kind  of 
analysis.  The  decomposition  of  steam  by  hot  iron,  with 
the  formation  of  magnetic  oxide  of  iron  and  hydrogen 
(p.  19)  is  also  an  analysis,  when  the  composition  of  the  oxide 
of  iron  formed  has  been  established  as  the  same  composition 
as  that  of  the  oxide  of  iron  formed  by  burning  iron  in 
oxygen. 

The  composition  of  a  substance  is  determined  by  synthe- 
sis when  the  elements  of  which  it  is  composed  are  put  to- 
gether to  form  it,  as  when  hydrogen  is  burned  in  oxygen, 
forming  water.  A  synthesis  of  water  may  also  be  made  by 
passing  hydrogen  over  heated  copper  oxide,  giving  copper 
and  water.  An  analysis  or  a  synthesis  may  be  qualitative, 
giving  simply  the  elements  of  which  the  substance  is  com- 
posed, or  quantitative,  giving  the  proportion  of  each  element 
present. 

Electrolysis  of  Sulfuric  Acid. — A  very  simple  and  a 
roughly  quantitative  analysis  of  water  can  be  made  by 
passing  an  electric  current  through  dilute  sulfuric  acid, 

42 


ELECTROLYSIS 


using  the  apparatus  shown  in  Fig.  18.  Electrodes  of  plat- 
inum in  each  arm  of  the  U-tube  are  connected  with  the 
poles  of  an  electric  battery  by  means  of  platinum  wires 
which  pass  through  the  glass.  It  can  be  shown  that  the 
hydrogen  of  the  sulfuric  acid  is  carried  toward  the  negative 
electrode  through  the  solution,  while  the  rest  of  the  acid, 
composed  of  sulfur  and  oxygen,  is  carried 
toward  the  positive  electrode.  At  the 
surface  of  the  negative  electrode  the 
hydrogen  is  liberated  as  a  gas,  while  at 
the  positive  electrode  oxygen  is  liber- 
ated. The  volume  of  the  hydrogen  is 
twice  that  of  the  oxygen. 

Since  a  liter  of  hydrogen  weighs  0.09 
gram,  and  a  liter  of  oxygen  1.429  grams, 
the  proportion  by  weight  is  0.09  X 
2  :  1.429  =  1  :  7.94,  approximately. 
The  experiment  is  not  suitable,  how- 
ever, for  an  accurate  determination  of 
the  composition  of  water. 

As  the  hydrogen  and  oxygen  are  liber- 
ated in  the  same  proportion  in  which 
they  are  found  in  water  and  as  the  total 
amount  of  sulfuric  acid  in  the  solution 
remains  unchanged  the  experiment  is 
often  spoken  of  as  the  decomposition  of  water  by  elec- 
tricity. In  consideration  of  the  final  result  this  is  correct, 
but  it  must  not  be  overlooked  that  the  motion  of  the 
two  parts  of  the  sulfuric  acid,  hydrogen  in  one  direction 
and  the  sulfur  and  oxygen  together  in  the  other,  arc  an 
essential  part  of  the  process  of  electrolysis. 

The  decomposition  of  a  substance  by  passing  a  current  of 
electricity  through  it  is  called  electrolysis.  The  poles  con- 
nected with  the  battery  or  dynamo  which  furnishes  the 
current  of  electricity  are  called  electrodes.  The  substance 


FIG.  18. 


44 


COMPOSITION  OF  WATER 


team 


-18 


=20 


which  is  decomposed  is  called  an  electrolyte.  The  positive, 
electrode,  toward  which  the  negative  constituent  of  the 
electrolyte  moves,  is  called  the  anode.  The  negative  elec- 
trode, toward  which  the  positive  constituent  of  the  elec- 
trolyte moves  is  called  the  cathode.  The  hydrogen  of  the 
sulfuric  acid  is  positive  and  moves  toward  the 
cathode,  while  the  sulfur  and  oxygen  together  are 
negative  and  move  toward  the  anode. 

Volumetric  Composition  of 
Water.  —  The  proportion  in 
which  hydrogen  and  oxygen 
combine  by  volume  to  form 
water  may  be  more  accurately 
determined  by  mixing  measured 
quantities  of  oxygen  and  hydro- 
gen in  a  graduated  tube,  ex- 
ploding the  mixture  by  means  of 
an  electrical  spark  passed  be- 
tween platinum  wires  sealed 
through  the  walls  of  the  tube 
and  measuring  the  volume  of 
oxygen  or  hydrogen  remaining 
uncombined  after  the  explosion. 
Thus  if  we  introduce  11  cc.  of| 
oxygen  and  25  cc.  of  hydrogen 
into  the  tube  shown  in  Fig.  19' 
and  then  explode  the  mixture, 
it  will  be  found  that  3  cc.  of  hydrogen  will  remain  uncom- 
bined. Or  if  15  cc.  of  oxygen  are  mixed  with  18  cc.  of 
hydrogen,  after  explosion  6  cc.  of  oxygen  will  remain.  The 
gases  must,  of  course,  be  measured  at  the  same  tempera- 
ture and  pressure  in  each  case.  If  the  tube  is  heated  and 
the  water  formed  by  the  combination  converted  into 
steam  (Fig.  20),  it  is  found  that  after  taking  account  of  the 
increase  in  temperature  and  correcting  the  volume  of  the 


FIG.    19. 


DUMAS'S  EXPERIMENT 


45 


gas  back  to  the  original  temperature  there  will  be  25  cc. 
of  hydrogen  and  steam  in  the  first  case  and  24  cc.  of  oxygen 
and  steam  in  the  second  case.  A  little  study  of  these  re- 
sults will  show  that  one  volume  of  oxygen  combines  with 
two  volumes  of  hydrogen  to  give  two  volumes  of  steam. 
This  may  be  expressed  graphically  by  the  following  diagram : 


Oxygen 


Hydrogen 

Hydrogen 

Steam 

Steam 

Composition  of  Water  by  Weight.— The  composition  of 
water  by  weight  has  been  determined  by  reducing  a  weighed 
amount  of  copper  oxide  to  metallic  copper  by  means  of 
hydrogen  and  collecting  and  weighing  the  water  formed. 
Thus  hydrogen  from  the  generator,  F(  Fig.  21),  is  passed 
first  through  a  series  of  tubes  to  purify  and  dry  the  gas  and 
then  through  the  heated  bulb,  B,  containing  copper  oxide, 
where  the  hydrogen  takes  oxygen  away  from  the  oxide 
and  combines  with  it  to  form  water.  The  water  is  collected 
partly  in  a  bulb,  BI,  just  beyond  the  one  containing  copper 
oxide,  partly  in  tubes  containing  calcium  chloride  or  other 
substances  which  absorb  vapor  of  the  water,  which  might 
otherwise  escape.  The  loss  in  weight  of  the  copper  oxide 
gives  the  weight  of  the  oxygen,  and  the  gain  in  weight  of 
the  bulb  and  tubes  in  which  the1  water  is  collected  gives 
the  weight  of  the  water  formed.  With  this  apparatus  the 
hydrogen  cannot  be  weighed  directly,  but  its  weight  is 
found  by  subtracting  the  weight  of  the  oxygen  from  the 
weight  of  the  water.  Dumas,  the  French  chemist  who 
first  used  this  method  very  carefully,  did  not  succeed  in 
obtaining  very  accurate  results  with  it. 

By  means  of  the  apparatus  shown  in  Fig.  22  it  is  possible 
to  weigh  the  hydrogen  directly.  The  bulb  A  is  filled  with 
copper  oxide  and,  after  exhausting  it  of  air  and  closing  the 


46 


COMPOSITION  OF  WATER 


RATIO  OF  OXY.GEN  TO  HYDROGEN 


47 


stopcock,  it  is  weighed.  It  is  then  placed  in  an  air-bath  so 
that  it  can  be  heated,  while  the  side  tube  B  is  cooled  and 
the  tube  C  is  connected  with  an  apparatus  giving  pure,  dry 
hydrogen.  As  this  passes  over  the  hot  copper  oxide  it  is 
oxidized  to  water  and  the  latter  is  condensed  in  the  tube  B. 
After  a  considerable  amount  of  water  has  collected,  the 
apparatus  is  cooled  and  weighed  again.  Evidently  the 
gain  in  weight  is  the*  weight  of  the  hydrogen  which  has  en- 
tered the  apparatus,  as  all  of  the  oxygen  of  the  copper  oxide 


B 


FIG.    22. 


remains  within  the  apparatus  either  as  unreduced  copper 
oxide  or  as  water.  By  heating  the  apparatus  and  side  tube 
after  connecting  it  with  another  bulb  from  which  the  air 
has  been  exhausted,  the  water  can  be  driven  out  and 
collected.  The  difference  between  the  weight  of  the  bulb  at 
first  and  the  weight  after  removing  the  water  gives  the 
weight  of  the  oxygen  which  has  combined  with  the  hydrogen. 
The  apparatus  in  which  the  water  is  collected  may  also  be 
weighed  and  the  weight  of  the  water  formed  determined. 
The  results  of  experiments  by  this  method  showed  that  the 
ratio  between  the  weights  of  hydrogen  and  oxygen  in  water 
is  H  :  O  =  1  :  7.938.  Determinations  by  other  still  more 
accurate  methods  have  given  almost  exactly  the  same 
result. 


48  COMPOSITION  OF  WATER 

Law  of  Constant  Proportion. — The  more  carefully  the 
experiments  for  determining  the  composition  of  water  are 
carried  out  the  more  exactly  do  the  results  agree  with  the 
value  for  the  ratio  which  has  been  stated  above.  This 
illustrates  the  law  of  constant  proportion,  which  has  already 
been  given  (p.  3). 

Law  of  Combining  Proportions. — We  may  select  for  each 
element  some  number  which  may  always  be  used  to  repre- 
sent the  proportion  of  the  element  which  enters  into  combi- 
nation with  other  elements.  This  may  be  illustrated  by 
the  following  list  of  compounds:1 

Water  Cuprous  Oxide  Cupric  Sulfide        Hydrogen  Sulfide 

H  :  0          0  :  Cu  Cu  :  S  S  :  H 

1:8  8  :  63.6  63.6  :  32  32:2 

Hydrochloric  Cupric  Chloride  Cupric  Oxide         Chlorine 
Acid                                                                                 Dioxide 

H:C1                         Cl:Cu  Cu :  O               0 :  Cl 

2:71                         71:  63.6  63.6 :  16  16 :  17.75 

These  compounds  are  selected  so  that  each  contains  one 
element  of  the  preceding  and  one  of  the  following  compound. 
The  table  might  be  extended  to  contain  a  thousand  com- 
pounds and  it  would  be  found  that  in  all  compounds  of 
oxygen  the  amount  combining  with  any  other  element 
would  be  either  8  parts  or  8  parts  multiplied  or  divided  by 
some  'whole  number.  In  the  same  way  the  amount  of 
hydrogen  would  be  1,  2,  3,  4,  or  more  parts.  The  propor- 
tion of  copper  entering  into  combination  with  other  elements 
is  63.6  parts  for  the  cases  given,  and  if  the  table  were  ex- 
tended to  contain  other  compounds  of  copper  the  amount 
of  that  element  combining  with  other  elements  would 
always  be  63.6  parts  or  some  multiple  or  submultiple 
of  63.6  parts.  We  might  select  the  combining  proportions, 

1  For  the  sake  of  simplicity  round  numbers  are  used  here  and  elsewhere. 


LAW  OF  MULTIPLE  PROPORTIONS  49 

H  =  1,  0  =  8,  Cu  =  63.6,  S  =  32,  Cl  =  71,  and  we  could  then 
express  the  composition  of  every  compound  of  these  ele- 
ments by  these  numbers  or  by  multiples  or  submultiples 
of  these  numbers.  The  values  O  =  16  and  Cl  =  35.5  are 
practically  used  in  place  of  8  and  71. 

The  combining  proportions  which  have  been  selected  are 
called  atomic  weights,  because  it  is  believed  that  they  repre- 
sent the  relative  weights  of  those  smallest  particles  of  the 
elements,  which  are  called  atoms  (p.  29).  A  table  of 
atomic  weights  is  given  on  p.  164  and  on  the  inside  of  the 
first  cover. 

Hydrogen  Peroxide. — There  is  a  second  compound  of 
oxygen  and  hydrogen  which  contains  a  larger  proportion 
of  oxygen  than  water  does.  The  chemical  name  of  water  is 
hydrogen  oxide,  but  that  name  is  very  rarely  used.  Water 
contains,  in  round  numbers,  1  part  of  hydrogen  to  8  parts 
of  oxygen.  Hydrogen  peroxide  contains  1  part  of  hydrogen 
to  16  parts  of  oxygen  and  the  name  given  to  it  indicates 
this  composition,  the  prefix  per  meaning  "more  of."  In 
this  case  it  means  that  the  compound  contains  more  oxygen 
than  common  water. 

Hydrogen  peroxide,  when  pure,  has  a  much  higher 
specific  gravity  than  water;  It  is  a  colorless  liquid  and 
it  is  very  unstable.  When  pure  it  is  liable  to  decompose 
explosively  into  water  and  oxygen  gas. 

Nearly  all  explosions  depend  on  the  formation  of  a  large 
volume  of  gas  from  some  solid  or  liquid  substance. 

A  solution  of  hydrogen  peroxide  in  water  is  more  stable 
than  the  pure  compound  and  is  used  by  dentists  and  others 
as  a  germicide.  Hydrogen  peroxide  is  also  a  good  bleaching 
agent,  especially  for  hair  or  silk. 

Law  of  Multiple  Proportions. — The  weight  of  oxygen 
combined  with  1  part  by  weight  of  hydrogen  is  exactly 
twice  as  great  in  hydrogen  peroxide  as  it  is  in  water.  The 
relation  may  also  be  given  by  saying  that  there  is  twice  as 


50  COMPOSITION  OF  WATER 

much  hydrogen  in  water,  for  1  part  of  oxygen,  as  there 
is  in  hydrogen  peroxide.  Stated  in  the  form  of  a  proportion 
these  are: 

Water  Hydrogen  Peroxide 

H:  0  H  :  0 

1:8  1  :  16 

or  2  :  16  2  :  32 

or  K  :  1  He  :  1 

In  many  other  cases  two  elements  form  two  or  more 
compounds  with  each  other  and  in  every  case  relations 
similar  to  this  are  found.  This  has  given  us  the  law  of 
multiple  proportions:  //  two  elements  combine  in  different 
proportions  to  form  two  different  compounds,  when  we  con- 
sider a  fixed  amount  of  one  element,  the  amounts  of  the  other 
element  which  combine  with  this  amount  will  bear  a  simple 
ratio  to  each  other. 

The  following  still  more  striking  illustration  of  this  law 
may  be  given: 


N 

Nitrous  oxide 14 

Nitric  oxide 14 

Nitrogen  trioxide 14 

Nitrogen  tetroxide 14 

Nitrogen  pentoxide 14 


O  N:  O 

8  or  N20  28  :  16 

16  or  NO  28  :  32 

24  or  N203  28  :  48 

32  or  N204  28  :  64 

40  or  N205  28  :  80 


A  study  of  this  law  shows  that  it  follows,  necessarily, 
from  the  law  of  combining  proportions  and  that  the  law 
of  combining  proportions  is  more  general  in  its  application 
and  more  important.  The  law  of  multiple  proportions  is 
simpler,  however,  and  was  discovered  first. 

The  Atomic  Theory. — A  number  of  reasons,  have  been 
given  for  believing  that  gases  and  other  substances  consist 
of  very  small  particles,  which  are  called  molecules.  The 
law  of  multiple  proportions  was  discovered  by  Dalton 
about  1804  and  in  trying  to  find  some  reason  for  it  he  came 
to  the  conclusion  that  the  molecules  of  compounds  must  be 


SYMBOLS.     FORMULAS  51 

formed  by  the  union  of  still  smaller  particles  which  he 
called  atoms  of  the  elements.  He  supposed  further  that 
each  atom  of  an  element  weighs  the  same  as  every  other 
atom  of  the  same  element,  but  that  the  atoms  of  different 
elements  have  different  weights.  He  also '  supposed  that 
chemical  combination  always  takes  place  between  atoms. 
If  this  theory  is  true,  the  laws  of  constant  proportion,  of 
combining  weights,  and  of  multiple  proportions  follow 
directly  from  it.  Further  than  that,  while  we  cannot 
weigh  an  atom  or  molecule  by  any  ordinary  method,  we 
can  determine  the  relative  weights  of  the  atoms  by  deter- 
mining the  composition  by  weight  of  the  compounds  which 
they  form.  Thus  if  we  can  discover  in  some  way  (see  p.  131) 
that  a  molecule  of  water  consists  of  two  atoms  of  hydrogen 
combined  with  one  atom  of  oxygen,  when  we  have  deter- 
mined that  1  part  by  weight  of  hydrogen  combines  with 
8  parts  by  weight  of  oxygen  to  form  water  we  can  at  once 
say  that  an  atom  of  oxygen  is  16  tunes  as  heavy  as  an  atom 
of  hydrogen.1  Evidently  the  atomic  weights,  if  we  can 
determine  them  by  such  a  process,  will  be  the  most  satis- 
factory numbers  to  use  as  the  combining  proportions  of  the 
elements,  to  express  the  composition  of  compounds. 

Symbols,  Formulas. — It  has  been  found  convenient  to  use 
the  symbol  of  an  element  to  stand  for  an  atom  of  the 
element.  The  number  of  atoms  of  each  element  in  a  mole- 
cule of  a  compound  is  designated  by  a  small  figure  placed 
below  the  line  and  to  the  right  of  each  symbol.  Thus  the 
composition  of  a  molecule  of  water  is  represented  by  the 
formula  H2O,  which  means  that  one  molecule  of  water 
contains  two  atoms  of  hydrogen  and  one  atom  of  oxygen. 
The  formula  might  be  a  little  clearer,  perhaps,  if  it  were 
written  H20i,  but  the  number  one  is  always  understood 

1  Round  numbers  are  used.  The  true  value  is  15.88.  As  a  matter  of 
convenience,  however,  16  has  been  selected  as  the  atomic  weight  of  oxygen 
and  the  accurate  atomic  weight  of  hydrogen  is  1.0077. 


52  COMPOSITION  OF  WATER 

when  no  figure  is  used  with  a  symbol.  The  formula  of 
hydrogen  peroxide  is  H2O2,  that  of  sulfur  dioxide  SO2, 
that  of  carbon  dioxide  CO2,  of  phosphorus  pentoxide 
P2O5,  of  magnetic  oxide  of  iron  Fe3O4,  of  sodium  hydroxide 
NaOH. 

The  atomic  weights,  in  round  numbers,  for  the  elements 
most  used  in  these  formulas  are: 

H  =  1,  O  =  16,  S  =  32,  C  =  12,  P  =  31, 

Fe  =  56,  Na  =  23. 

Distinction  between  Parts  by  Weight  and  Atoms. — If 
we  remember  the  atomic  weight,  the  formula  of  a  compound 
tells  us  its  composition  by  weight.  Water,  H2O,  contains 
two  atoms  of  hydrogen  for  one  atom  of  oxygen  and  it 
contains  2  parts  by  weight  of  hydrogen  for  16  parts 
by  weight  of  oxygen.  It  should  be  noticed,  especially, 
that  the  formula  does  not  mean  that  water  contains  2  parts  of 
hydrogen  for  1  part  of  oxygen,  though  such  a  misinterpreta- 
tion of  the  formula  is  often  made' by  beginners.  Phosphorus 
pentoxide,  P2O5,  contains  twice  31,  or  62  parts  of  phosphorus 
combined  with  five  times  16,  or  80  parts  of  oxygen.  Sodium 
hydroxide,  NaOH,  contains  23  parts  of  sodium,  16  parts 
of  oxygen  and  1  part  of  hydrogen. 

Equations. — Symbols  and  formulas  may  be  combined 
in  equations  which  furnish  a  very  concise  statement  of  what 
happens  in  chemical  reactions.  Thus  the  equation: 

Na  +  H2O  =  NaOH  +  H 

means  that  one  atom  of  sodium  acts  on  one  molecule  of 
water  to  form  one  molecule  of  sodium  hydroxide,  NaOH, 
and  one  atom  of  hydrogen.  It  means,  also,  that  23  parts 
of  sodium  with  18  parts  of  water  give  40  parts  of  sodium 
hydroxide  and  1  part  of  hydrogen. 
The  equation: 

3Fe  +  4O  =  Fe3O* 


EQUATIONS.     CALCULATIONS  53 

means  that  three  atoms  of  iron  combine  with  four  atoms 
of  oxygen  to  form  one  molecule  of  the  magnetic  oxide;  also 
that  3  X  56  =  168  parts  of  iron  combine  with  4  X  16  =  64 
parts  of  oxygen  to  form  168  +  64  =  232  parts  of  the 
magnetic  oxide  of  iron. 

Writing  Equations. — Students  often  make  the  mistake 
of  memorizing  equations.  This  ought  never  to  be  done. 
The  first  step  should  always  be  to  learn  the  formulas  of  the 
compounds  which  act  on  each  other  and  the  formulas  of 
the  products  formed  in  the  reaction.  With  these  formulas 
as  a  starting  point,  the  equation  should  be  deduced  logically. 

Thus  when  steam,  H2O,  is  passed  over  heated  iron,  Fe, 
magnetic  oxide  of  iron,  Fe3O4,  and  hydrogen,  H,  are  formed. 
With  this  starting  point  we  write: 

Fe  +  H2O-»Fe304  +  H.,1 

On  looking  at  these  formulas  we  see  that  there  are  four 
atoms  of  oxygen  on  the  right  side  and  only  one  on  the 
left ;  there  are  also  three  atoms  of  iron  on  the  right  and  only 
one  on  the  left.  We  must  have,  therefore: 

3Fe  +  4H20-»Fe304  +  H2 

Looking  again  at  this  we  see  that  there  are  4  X  2  or  8 
atoms  of  hydrogen  on  the  left  and  we  must  have  the  same 
number  on  the  right.  The  equation  then  becomes: 

3Fe  +  4H2O  =  Fe3O4  +  4H2 

If  we  wish  to  express  the  fact  that  the  reaction  is  re- 
versible, we  may  use  two  arrows  in  place  of  the  sign  of 
equality.  Thus: 

3Fe  +  4H20<F±Fe3O4  +  4H2 

Calculations. — As  a  formula  indicates  the  composition 
of  a  compound  in  terms  of  the  atomic  weights  the  percentage 
composition  can  always  be  calculated  from  it.  The  formula 

1  The  reason  for  writing  the  formula  of  hydrogen  H?  and  that  of  oxygen 
O2  will  be  given  later  (p.  136). 


54  COMPOSITION  OF  WATER 

of   potassium    chlorate   is    KC103.     The    composition    by 

weight  is,  from  this  formula: 

Potassium,  K  =   39  parts 
Chlorine,  Cl     =   35.5  parts 
Oxygen,  03       =   3  X  16  or  48  parts 
Total  =   122.5  parts 

The  percentage  of  oxygen  must  be   100  „  X  100  =  39.2 


per  cent. 

If  we  have  the  equation  for  any  reaction  and  know  the 
amount  of  any  substance  used  or  formed  in  the  reaction 
we  can  calculate  the  amounts  of  every  other  substance 
involved. 

Suppose  we  have  25  grams  of  iron,  Fe,  and  wish  to  know 
how  much  water  it  will  decompose  and  what  weights  of 
magnetic  oxide  and  of  hydrogen  will  be  formed. 

The  first  step  in  solving  any  problem  of  this  sort  should 
be  to  write  the  equation  for  the  reaction  : 

3Fe  +  4H2O  =  Fe3O4  +  4H2 

Next  write  the  formulas  of  the  two  substances  to  be 
considered,  in  the  form  of  two  identical  ratios,  using  the 
number  which  is  used  before  each  formula  in  the  equation 
before  the  same  formula  in  the  ratios.  Beneath  the  for- 
mulas of  the  first  ratio  write  the  corresponding  atomic  or 
molecular  weights.  Beneath  the  formula  of  the  known 
substance  in  the  second  ratio  write  the  amount  of  that  sub- 
stance stated  in  the  problem.  Beneath  the  other  formula 
in  the  second  ratio  write  "x".  Finally  solve  the  proportion 
for  x. 

Thus  to  calculate  the  amount  of  water,  above: 

3Fe:4H2O  ::3Fe:4H20 
3  X  56:4  X  18::  25   :    x 
168:72::25    :    x 

x  =  —  =  10.71  grams  of  water 

.LOo 


SUMMARY.     WATER,  ATOMIC  THEORY  55 

In  a  similar  manner  the  amounts  of  magnetic  oxide  and 
of  hydrogen  may  be  calculated. 

The  student  should  also  notice  that  since  168  parts  of 
iron  will  decompose  72  parts  of  water,  the  amount  of  water 

is  found  by  multiplying  the  weight  of  iron,  25  grams,  by 
72 

Problems  of  this  sort  may  be  solved  by  using  a  frac- 


tion in  this  way  without  writing  the  formal  proportion. 
The  fraction  is  less  likely  to  be  used  mechanically  than  is  the 
proportion. 

SUMMARY 

The  composition  of  substances  is  determined  by  analysis, 
the  separation  into  the  elements  of  which  they  are  com- 
posed; or  by  synthesis,  preparation  by  putting  the  elements 
together. 

A  simple  analysis  of  water  may  be  made  by  the  electroly- 
sis of  dilute  sulfuric  acid,  giving  two  volumes  of  hydrogen 
to  one  of  oxygen. 

In  electrolysis  the  poles  of  the  battery  put  into  the  elec- 
trolyte are  called  electrodes. 

The  positive  electrode  is  the  anode,  the  negative  is  the 
cathode. 

A  mixture  of  measured  volumes  of  oxygen  and  hydrogen 
may  be  exploded,  and  by  measuring  the  volume  of  the  gas 
remaining  an  accurate  synthesis  of  water  can  be  made. 

By  passing  hydrogen  into  an  evacuated  apparatus  con- 
taining copper  oxide,  and  afterward  removing  the  water 
formed,  a  synthesis  of  water  by  weight  can  be  effected. 
The  gain  in  weight  gives  the  weight  of  the  hydrogen;  the 
loss  after  removing  the  water  gives  the  weight  of  the  oxygen. 
The  water  may  also  be  collected  and  weighed. 

A  pure  compound  always  has  exactly  the  same  composi- 
tion by  weight  (law  of  constant  proportion). 

A  number  may  be  selected  for  each  element  which  repre- 


56  COMPOSITION  OF  WATER 

sents  the  proportion  by  weight  in  which  the  element  enters 
into  combination  with  any  other  element  (law  of  combining 
proportions).  The  most  convenient  numbers  for  this  pur- 
pose are  the  atomic  weights  of  the  elements. 

Hydrogen  peroxide  is  a  second  compound  of  hydrogen  and 
oxygen.  It  contains  twice  as  much  oxygen  for  a  given 
weight  of  hydrogen  as  water  does. 

Hydrogen  peroxide  and  water  illustrate  the  law  that 
whenever  two  elements  form  two  compounds  having  a 
different  composition  the  amounts  of  one  which  combine 
with  a  fixed  amount  of  the  other  will  bear  a  simple  ratio 
to  each  other  (law  of  multiple  proportions). 

The  atomic  theory  gives  a  simple  and  satisfactory  expla- 
nation of  the  laws  of  constant  proportions,  combining  pro- 
portions and  multiple  proportions. 

Symbols  designate: 

(a)  An  element.     H  stands  for  hydrogen;  Fe  for  iron. 

(6)  An  atom  of  an  element.  0  stands  for  an  atom  of 
oxygen. 

(c)  A  gram  atom  of  the  element.  Fe  stands  for  56  grams 
of  iron. 

Formulas  designate: 

(a)  A  compound.  H2O  stands  for  water;  KC1O3  for 
potassium  chlorate. 

(6)  A  molecule  of  a  compound.  H2O2  stands  for  a  mole- 
cule of  hydrogen  peroxide;  O2  stands  for  a  molecule  of 
oxygen. 

(c)  A  gram-molecule  of  a  compound.     H2O  stands  for 
18  grams  of  water. 

(d)  The  composition  by  weight  of  a  compound.   H2O 
indicates  that  water  is  composed  of  2  parts  of  hydrogen 
combined  with  16  parts  of  oxygen. 

Equations  are  used  to  represent: 

(a)  The  substances  which  react  and  the  substances 
formed  in  a  chemical  reaction. 


EXERCISES.     WATER,  ATOMIC  THEORY  57 

(6)  The  atomic  composition  of  each  substance  involved. 
(c)  The  quantities  of  the  substances. 
Equations  should  be  true : 
(a)  Chemically. 
(6)  Algebraically. 

Formulas  and  equations  may  be  used  as  the  basis  for 
a  great  variety  of  calculations. 

EXERCISES 

1.  Write  equations  to  represent  the  burning  of  sulfur,  charcoal, 
phosphorus  and  iron.     The  formulas  of  the  compounds  formed 
are  S02,  C02,  P205  and  Fe304. 

2.  Write  the  equations  for  the  decomposition  of  oxide  of  mer- 
cury, HgO,  and  of  potassium  chlorate,  KClOs. 

3.  At  a  high  temperature  manganese  dioxide,  Mn02,  may  be 
decomposed  into  oxygen,  02,  and  another  oxide,  Mn304.     Write 
the  equation  for  the  reaction. 

4.  Write  the  equations  for  the  action  of  hydrochloric  acid,  HC1, 
and  of  sulfuric  acid,  H2S04,  on  zinc  and  on  iron.     The  compounds 
formed   are   zinc   chloride,    ZnCl2,   ferrous   chloride,  FeCU,  zinc 
sulfate,  ZnS04,  and  ferrous  sulfate,  FeS04. 

5.  Using    the    atomic    weights,    H  =  1,  0.  =  16,  Cu  =  63.6, 
S  =  32,  Cl  =  35.5,  write  the  formulas  used  in  illustrating  the  law 
of  combining  proportions. 

6.  Calculate  the  percentage  composition  of  mercuric  oxide,  HgO 
(Hg  =  200),    potassium    chlorate,    KC103  (K  =  39,  Cl  =  35.5), 
and  sulfuric  acid,   H2SO4  (S  =  32). 

7.  If  a  gasometer  holds  20  liters,  how  much  potassium  chlorate 
will  be  required  to  give  enough  oxygen  to  fill  it?     One  liter  of 
oxygen  weighs  1.429  grams. 

8.  How  much  zinc  and  how  much  sulfuric  acid  will  be  required 
to  fill  the  same  gasometer  with  hydrogen?     One  liter  of  hydrogen 
weighs  0.09  gram. 

9.  Air  contains  21  per  cent  by  volume  of  oxygen.     How  many 
liters  of  air  will  be  required  to  burn  one  pound  (453  grams)  of 
sulfur  (S  =  32)?     How  many  liters  will  be  required  to  burn  one 
pound  of  charcoal,  supposing  it  to  be  pure  carbon  (C  =  12)? 


CHAPTER  VI 
PROPERTIES  AND  USES  OF  WATER.     VAPOR  PRESSURE 

Freezing  Water. — When  ice  is  heated  it  begins  to  melt 
at  0°  and  if  the  ice  and  the  water  formed  (solid  and  liquid 
phases)  are  kept  in  intimate  contact  the  temperature  of 
the  mixture  will  remain  at  0°  till  all  the  ice  is  melted.  If 
water  containing  no  ice  is  cooled,  its  temperature  may  fall 
considerably  below  the  freezing  point  before  any  ice  begins 
to  form.  Water  in  this  condition  is  said  to  be  supercooled. 
It  is  in  an  unstable  condition  and  if  some  ice  is  thrown  into 
this  supercooled  water  part  of  the  water  in  contact  with 
the  ice  will  freeze  and  in  freezing  will  give  out  heat  to  the 
rest  of  the  water  until  the  temperature  of  the  whole  rises 
to  zero.  Ice  and  water  cannot  be  in  contact  with  each 
other  at  any  other  temperature. 

If  two  dishes,  one  containing  a  kilogram  of  ice  and  the 
other  containing  a  kilogram  of  water  at  the  freezing  point, 
are  placed  over  two  gas  burners  in  such  a  manner  that  each 
receives  heat  at  the  same  rate,  when  the  ice  is  just  melted 
and  the  water  in  contact  with  it  is  still  at  zero  the  water 
in  the  other  dish  will  be  found  to  have  reached  a  tempera- 
ture of  80°.  In  other  words,  it  requires  as  much  heat  to 
melt  a  pound  of  ice  without  changing  its  temperature  as  it 
does  to  raise  the  temperature  of  a  pound  of  water  80°  or 
to  raise  the  temperature  of  80  pounds  of  water  one  degree. 
The  latent  heat  of  fusion  of  ice  is  said  to  be  80  calories 
(accurately  79.63  calories). 

A  calorie  is  the  heat  required  to  raise  the  temperature  of 
one  kilogram  of  water  one  degree.  The  small  calorie,  the 

58 


CHANGES  OF  VOLUME  OF  WATER  59 

heat  required  to  raise  the  temperature  of  a  gram  of  water 
one  degree,  is  also  often  used. 

Boiling  of  Water. — If  water  is  heated  to  100°  under  at- 
mospheric pressure  (760  mm.),  bubbles  of  steam  will  form 
and  rise  through  the  water,  and  if  some  means  is  provided  to 
secure  the  continual  formation  of  minute  bubbles  in  the 
water,  the  water  will  continue  to  boil  and  the  temperature 
will  remain  at  100°  till  all  of  the  water  has  been  converted 
into  steam.  If  no  provision  for  starting  bubbles  of  steam 
is  made,  the  water  may  become  superheated  to  a  temperature 
considerably  above  100°,  but  if  bubbles  of  gas  are  introduced 
in  superheated  water  the  water  will  boil  rapidly  and  the 
temperature  will  fall  to  100°,  the  true  boiling  point. 

If'  two  flasks,  one  containing  a  kilogram  of  water  at  0° 
and  the-  other  containing  a  kilogram  of  water  at  100°,  are 
placed  over  two  burners  so  that  the  same  quantity  of 
heat  enters  each,  the  water  at  100°  will  boil  away  without 
changing  its  temperature,  but  when  the  temperature  of  the 
water  in  the  first  flask  has  risen  to  100d  less  than  one-fifth 
of  the  water  in  the  second  flask  will  have  been  changed  to 
steam.  If  another  flask  at  0°  is  placed  over  the  first  burner 
and  left  till  its  temperature  reaches  100°  and  this  is  repeated 
until  all  the  water  in  the  second  flask  is  converted  into 
steam,  it  will  be  found  that  the  water  in  five  flasks  has  been 
heated  from  0°  to  100°  and  that  the  water  in  the  sixth 
flask  will  be  at  a  temperature  of  36°.  The  latent  heat  of 
vaporization  of  water  at  100°  is  said  to  be  536  calories  (accu- 
rately 536.6).' 

Changes  of  Volume  of  Water  for  Changes  of  Tempera- 
ture.— With  very  few  exceptions,  substances  of  all  kinds, 
whether  solids,  liquids  or  gases,  expand  when  heated  and 
contract  when  cooled.  A  number  of  liquids  expand  on 
solidifying,  although  few,  if  any,  expand  as  much  as  water 
when  it  freezes.  But  water  is  almost,  or  quite,  alone  in 
expanding  on  cooling  while  the  temperature  is  still  several 


60 


PROPERTIES  AND  USES  OF  WATER 


degrees  above  the  freezing  point.  When  water  is  cooled  it 
contracts  as  other  substances  do  till  a  temperature  of 
4°  above  the  freezing  point  is  reached.  If  cooled  further, 
it  begins  to  expand  and  8000  cc.  of  water  at  4°  will  fill  a 
volume  of  8001  cc.  at  0°.  On  freezing,  the  same  amount  of 
water  will  give  8727  cc.  of  ice.  The  expansion  of  water 
on  freezing  is  familiar  in  the  bursting  of  a  pitcher  or  a 
bottle  filled  with  water,  ink  or  other  aqueous  liquid,  when 
the  liquid  freezes.  While  the  increase  of  volume  from  4° 
to  0°  is  small,  it  is  very  important  practically,  since  this 
property  of  water  protects  the  ocean  and  other  large 
bodies  of  water  from  cooling  to  the  freezing 
point,  because  the  lighter,  cooler  water 
floats  on  the  surface  and  is  a  poor  con- 
ductor of  heat.  As  long  as  the  tempera- 
ture of  a  body  of  water  is  above  4°  the 
colder  water  on  top  falls  and  is  replaced 
by  .warmer  water  from  below,  but  at  4° 
this  method  of  cooling  (called  cooling  by 
convection)  ceases.  The  ice  which  finally 
forms  on  the  surface  still  further  protects 
the  mass  of  water  below  from  cooling  below 
a  temperature  of  4°,  because  it  is  a  very 
poor  conductor  of  heat  and  cannot  be 
cooled  by  convection. 

Vapor  Pressure  of  Water. — If  a  drop  of 
water  is  introduced  into  the  space  above 
the  mercury  in  a  barometer  (Fig.  23),  a  little 
of  the  water  will  evaporate  and  the  vapor 
of  the  water  will  exert  a  pressure  on  the  mercury  and  cause 
the  mercury  to  fall  about  20  mm.  (^5  inch)  if  the  tem- 
perature is  22°  (72°  F.) .  If  the  temperature  is  higher  the 
mercury  will  fall  further;  if  it  is  lower  it  will  fall  less.  For 
any  given  temperature  the  pressure  of  the  vapor  is  always 
the  same,  provided  only  water  and  vapor  are  present  in 


FIG.    23. 


YAPOR  PRESSURE 


61 


the  space  above  the  mercury.  If  a  few  drops  of  water  are 
introduced  into  air  which  has  been  carefully  dried  and 
which  is  contained  in  a  closed  bottle  (Fig.  24),  some  of  the 
water  will  evaporate  and  the  vapor  of  the  water  will  diffuse 
into  the  air  of  the  bottle  and  will  increase  the  pressure 
within  the  bottle,  as  shown  by  the  manometer  (A)  con- 
nected with  it.  This  increase  in  pressure  will  be  found  to  be 
the  same  as  the  pressure  of  the 
vapor  in  the  barometer  tube 
above  the  mercury.  This  is  be- 
cause the  molecules  of  water  in 
the  form  of  vapor,  which  is 
merely  a  liquid  or  solid  con- 
verted into  a  gas,  shoot  in 
between  the  molecules  of  air 
(p.  31)  and  ultimately  there  are 
just  as  many  molecules  of  vapor 
as  there  would  be  if  the  air  were 
not  present. 

In  any  mixture  of  gases  each 
gas  exerts  the  same  pressure  as 
it  would  exert  if  it  were  present 
alone  in  the  same  space  (Henry's 
law). 

If  the  tube  containing  the 
mercury  and  water  is  heated  to  100°  while  the  pressure  of  the 
air  is  760  mm.,  the  mercury  within  and  without  the  tube 
will  be  at  the  same  level.  In  other  words,  the  boiling  point 
of  water  is  that  temperature  at  which  the  pressure  of  the 
vapor  of  water  is  equal  to  the  pressure  of  the  atmosphere. 

Natural  Waters.— Rain  water  which  falls  in  the  open 
country,  after  it  has  been  raining  some  time  so  that  the 
substances  which  would  contaminate  the  water  have  been 
washed  out  of  the  air,  is  very  nearly  pure.  As  soon  as  the 
water  reaches  the  earth,  however,  it  begins  to  take  up 


FIG.    24. 


62  PROPERTIES  AND  USES  OF  WATER 

various  substances.  Carbonic  acid,  H2CO3,  formed  by 
the  oxidation  of  vegetable  and  animal  substances  ("  organic 
matter")  in  the  soil,  is  taken  up  by  the  water,  and  if  the  soil 
also  contains  small  particles  of  limestone,  or  calcium 
carbonate,  CaCO3,  as  is  usually  the  c&se,  this  dissolves  in 
the  carbonic  acid  of  the  water  and  makes  it  "hard." 
Such  a  water  requires  more  soap  than  soft  water  to  produce 
a  lather  because  the  soap  is  at  first  used  up  in  producing 
a  curdy  precipitate  with  the  calcium  of  the  calcium  car- 
bonate. On  boiling  the  water  the  carbonic  acid,  H2CO3, 
which  holds  the  calcium  carbonate  in  solution,  breaks  up 
into  carbon  dioxide,  CO2,  and  water  and  the  calcium  car- 
bonate is  precipitated.  This  causes  a  scale  in  teakettles 
and  boilers. 

Equilibrium  between  a  Liquid  and  its  Vapor.  Kinetic 
Theory. — In  discussing  the  kinetic  theory  of  gases  (p.  31) 
it  was  pointed  out  that  while  the  molecules  of  gases  are  so 
far  apart  and  in  such  rapid  motion  that  they  exert  little 
attraction  upon  each  other,  the  molecules  of  a  liquid  are 
held  together  by  attractive  forces.  It  does  not  follow  from 
this  that  the  molecules  of  a  liquid  are  not  also  in  motion. 
Indeed  we  cannot  well  explain  the  escape  of  a  vapor  from  the 
surface  of  a  liquid  without  supposing  that  the  molecules 
of  a  liquid  are  in  motion  and  that  occasionally  the  motion 
of  an  individual  molecule  is  sufficiently  rapid  so  that  it 
tears  itself  away  from  its  fellows  and  shoots  out  into  the 
space  above.  Such  a  molecule  then  becomes  a  gaseous 
molecule  and  it  will  continue  straight  on  until  it  meets  with 
some  obstacle.  On  the  other  hand,  an  occasional  molecule 
of  the  vapor  will  be  moving  toward  the  liquid  surface  and 
when  it  reaches  this  it  may  be  held  by  the  attraction  of 
the  molecules  of  the  liquid  and  become  a  part  of  the  liquid 
once  more.  At  any  given  temperature  the  number  of 
molecules  leaving  the  liquid  and  the  number  returning  to 
it  must  be  the  same,  when  a  state  of  equilibrium  is  reached. 


TABLE  OF  VAPOR  PRESSURE 


63 


If  the  volume  of  the  vapor  above  the  liquid  is  increased 
by  enlarging  the  space  filled  by  it,  as,  for  instance,  by  rais- 
ing the  piston  in  Fig.  25,  the  pressure  of  the  vapor  will  be 
decreased  in  accordance  with  the  laws  of  gases.  Immedi- 


VAPOR  PRESSURE  OF  ICE  AND  WATER 


TEMPERA- 
TURES 

PRESSURE  IN 
MILLIMETERS 
OP  MERCURY 

TEMPERA- 
TURES 

PRESSURE 

IN    MM. 

TEMPERA- 
TURES 

PRESSURE 

IN    MM. 

-10° 

2.0 

27° 

26.5 

100.5° 

773.7 

-  5° 

3.0 

•28° 

28.1 

101.0° 

787.6 

-  2° 

3.9 

29° 

29.8 

-   1° 

4.2 

30° 

31.6 

±    0° 

4.6 

31° 

33.4 

1° 

4.9 

32° 

35.4 

2° 
3° 

5.3 

5.7 

33° 
34° 

37.4 
39.6 

PRESSURE  IN 

4° 

6.1 

35° 

41.9 

ATMOSPHERES 

5° 

6.5 

40° 

55.0 

6° 

7.0 

50° 

92.2 

111.7° 

1.5 

7° 

7.5 

60° 

149.2 

120.6° 

2 

8° 

8.0 

70° 

233.8 

127.8° 

2.5 

9° 

8.6 

80° 

355.5 

133.9° 

3 

10° 

9.2 

90° 

526.0 

144.0° 

4 

11° 

9.8 

95° 

634.0 

159.2° 

6 

12° 

10.5 

96° 

657.7 

170.8° 

8 

13° 

11.2 

97° 

682.1 

180  .  3° 

10 

14° 

11.9 

98° 

707.3 

188  .  4° 

12 

15° 

12.7 

99° 

733.2 

195  .  5° 

14 

16° 

13.6 

99.1° 

735.9 

201.9° 

16 

17° 

14.5 

99.2° 

738.5 

207.7° 

18 

18° 

15.4 

99.3° 

741.2 

213.0° 

20 

19° 

16.4 

99.4° 

743.9 

224.7° 

25 

20° 

17.4 

99.5° 

746.5 

21° 

18.5 

99.6° 

749.2 

22° 

19.7 

99.7° 

751.  9* 

23° 

20.9 

99.8° 

754.6 

24° 

22.2 

99.9° 

757.3 

25° 

23.5 

100.0° 

760.0 

26° 

25.0 

100.1° 

762.7 

100.2° 

765.5 

100.3° 

768.2 

100  .  4° 

770.9 

64  PROPERTIES  AND  USES  OF  WATER 

ately,  however,  a  larger  number  of  molecules  will  leave 
the  surface  of  the  liquid  than  will  return  to  it  because 
fewer  strike  the  surface,  and  this  will  continue  till  the 
pressure  is  the  same  as  before.  If  the  piston  is  forced 
down  and  the  pressure  of  the  vapor  is  temporarily  increased, 
the  opposite  process  will  result.  For  each  temperature 
every  liquid  has  a  definite  pressure  for  its  vapor  in 
contact  with  the  liquid.  The  same  law  applies  to 
solids  which  are  volatile.  Ice  has  a  vapor  pressure 
which  decreases  as  it  grows  colder. 

The  table  on  page  63  gives  the  vapor  pressures 
of  ice  and  water  at  different  temperatures. 

Germs  of  Diseases  Carried  by  Water. — It  often 
happens  that  by  means  of  sewage  disease  germs 
coming  from  cases  of  typhoid  fever  and  from  other 
diseases  find  their  way  into  rivers  or  lakes  or  wells. 
Persons  who  drink  water  that  has  been  contam- 
inated in  this  manner  are  liable  to  contract  the 
disease  which  is  caused  by  these  germs.  The 
FIG  25  ^ac^er^a  which  are  the  exclusive  cause  of  such 
diseases  as  typhoid  fever  and  cholera  may  be  re- 
moved from  water  by  proper  filtration  or  they  may  be 
killed  by  boiling  the  water.  A  number  of  years  ago  there 
was  a  serious  epidemic  of  cholera  in  Hamburg.  People 
living  in  Hamburg  used  water  taken  directly  from  the  river 
Elbe.  The  people  living  in  Altona,  a  city  practically  con- 
tinuous with  Hamburg,  used  water  taken  from  the  same 
river  but  filtered  it  before  it  was  pumped  into  the  city  mains. 
Many  people  in  Hamburg  died  from  the  cholera  while  there 
were  very  few  cases  of  cholera  in  Altona.  In  very  many 
cases  epidemics  of  typhoid  fever  have  been  traced  directly 
to  impure  drinking  waters. 

Solutions. — If  salt  or  sugar  is  mixed  with  water,  a  certain 
amount  is  dissolved.  Substances  which  are  not  volatile 
lower  the  freezing  point  and  raise  the  boiling  point  of  water 


CHEMICAL  ACTIVITY.     IONIZATION  65 

or  .any  other  solvent  in  which  they  are  dissolved.  The 
lowering  of  the  freezing  point  is  familiar  in  the  freezing 
mixtures  of  ice  and  salt,  which  are  used  in  making  ice 
cream. 

Many  substances  are  much  more  easily  soluble  in  hot 
water  than  in  cold  water.  If  a  concentrated  solution  of  such 
a  substance  is  cooled  without  the  presence  of  any  of  the  solid 
substance,  a  supersaturated  solution  may  be  obtained,  but 
the  introduction  of  a  crystal  of  the  substance  will  cause 
a  part  of  it  to  separate  from  the  solution  exactly  as  crystals 
of  ice  will  cause  a  part  of  supercooled  water  to  freeze.  A 
supersaturated  solution  is  unstable  and  cannot  exist  when 
the  solution  and  crystals  of  the  substance  are  in  intimate 
contact.  A  solution  obtained  by  shaking  water  or  some 
other  solvent  with  a  solid  which  dissolves  in  it  till  no  more 
is  taken  up  is  said  to  be  saturated. 

Chemical  Activity  in  Solutions.  lonization. — When  a 
solution  of  salt,  NaCl,  is  added  to  a  solution  of  silver  nitrate, 
AgNO3,  a  white,  curdy  precipitate  of  silver  chloride,  AgCl, 
is  formed.  Similar  reactions  will  take  place  between  dis- 
solved substances  in  many  other  cases.  Such  actions 
occur  for  two  reasons:  first,  because  in  solution  the  sub- 
stances are  brought  into  intimate  contact;  second,  because  in 
solution  such  substances  as  sodium  chloride  and  silver 
nitrate  separate  more  or  less  completely  into  ions,  the 
sodium  chloride,  NaCl,  into  sodium,  Na+,  and  chloride, 
Cl~,  ions;  the  silver  nitrate  into  silver,  Ag+,  and  nitrate, 
N03~,  ions,  and  any  positive  ion  in  a  solution  may  combine 
with  any  negative  ion  in  the  same  solution. 

Electrolysis. — In  considering  the  decomposition  of  water 
by  electricity  (p.  43)  it  was  stated  that  under  the  influence 
of  the  electrical  current  the  hydrogen  of  the  sulfuric  acid, 
H2SO4,  moves  through  the  solution  in  one  direction  while 
the  sulfur  and  oxygen  move  together  in  the  opposite  direc- 
tion. This  is  because  the  sulfuric  acid  separates  into 

5 


66 


PROPERTIES  AND  USES  OF  WATER 


positive  hydrogen  ions,  H+,  and  negative  sulf ate  ions,  SO  j=. 
The  positive  ions  are  attracted  toward  the  cathode,  where 
they  are  discharged  and  liberated  as  free  hydrogen,  H2. 
The  sulfate  ions,  SO4=,  are  attracted  toward  the  anode, 
but  in  contact  with  the  anode  there  are  also  a  few  hydroxyl, 
OH~,  and  perhaps  oxygen,  0=,  ions,  which  are  more  easily 
discharged  than  the  sulfate  ions.  The  hydroxyl  or  oxygen 
ions  are  discharged,  therefore,  instead  of  the  sulfate  ions 
and  in  this  way  free  oxygen  is  liberated. 

40H  =  2H2O  +  O2 

Migration  of  Ions. — The  light  hydrogen  ions  pass  through 
the  solution  toward  the  cathode  much  more  rapidly  than 
the  sulfate  ions  pass  toward  the  anode.  This  different 
rate  of  transfer  for  different  ions  can  be  demonstrated 
by  using  the  apparatus  illustrated  in  Fig.  26.  The  U-tube 
of  the  apparatus  is  filled  with  a  solution  containing  4 
milligram  molecules  (146  mg.)  of  hydrochloric  acid,  divided 
equally  between  the  two  arms. of  the  tube.  A  platinum 
cathode  and  silver  anode  are  used.  The  hydrogen  ions, 


The  solution  on  this  side 

contains : 
At  first, 
73  mg.  HC1  = 
2  milligram  atoms  of  H 
2  milligram  atoms  of  Cl 
At  end, 

66.5  mg.  HC1  = 
1.82    milligram    atoms 

of  H 
1.82    milligram    atoms 

of  Cl 
1  milligram  atom  of  H 

is  liberated 

Amount  of  hydrogen 
transferred  across  the 
line  CD  =  0.82  mg. 
atoms. 


Cathode 


4- 

V 

i 

m 

3 

= 

/ 

' 

V 

\^j  —  *- 

D 

FIG.  26. 

The  solutior  on  this  side 

contains : 
At  first, 
73  mg.  HC1  = 
2  milligram  atoms  of  H 
2  milligram  atoms  of  Cl 
At  end, 
43  mg.  HC1  = 
1.18    milligram    atoms 

of  H 
1.18    milligram    atoms 

of  Cl 
1  milligram  atom  of  Cl 

combines  with  anode 
Amount '  of  chlorine 
transferred  across  the 
line  CD  =  0.18  mg. 
atoms. 


MIGRATION  OF  IONS  67 

H+,  discharged  at  the  cathode  are  liberated  as  free  hydrogen, 
H2,  while  the  chloride  ions,  Cl~,  liberated  at  the  anode 
combine  with  the  silver  to  form  silvei  chloride. 

At  the  beginning  of  the  experiment  there  is  the  same 
amount  of  hydrochloric  acid  in  each  arm  of  the  tube,  2 
milligram  molecules.  If  the  current  is  passed  through  the 
solution  till  1  milligram  atom  of  hydrogen  has  been 
liberated  at  the  cathode  and  1  milligram  atom  of  chlorine 
has  combined  with  the  anode,  it  is  evident  that  if  the 
hydrogen  ions  move  through  the  solution  in  one  direction 
at  the  same  rate  that  the  chloride  ions  move  the  other 
way  there  would  be  1J^  milligram  molecules  of  hydro- 
chloric acid  in  each  arm  of  the  tube.  Instead  of  this  an 
analysis  of  the  solution  in  the  two  arms  will  show  that  the 
cathode  arm  contains  1.82  milligram  molecules  of  hydro- 
chloric acid  while  the  anode  arm  contains  only  1.18  milli- 
gram molecules.  It  is  evident  from  this  that  0.82  of  the 
milligram  atom  of  hydrogen  liberated  at  the  cathode  have 
been  replaced  by  hydrogen  transferred  from  the  anode  side 
while  only  0.18  of  the  chlorine  removed  at  the  anode  have 
been  replaced  by  chlorine  transferred  from  the  cathode 
side.  It  follows  that  the  hydrogen  ion  migrates  toward 
the  cathode  more  than  four  times  as  fast  as  the  chloride 
ion  migrates  toward  the  anode. 

The  independent  migration  of  the  ions  in  a  solution  of 
an  electrolyte  is  one  of  the  strong  reasons  for  believing  that 
electrolytes  actually  separate  into  ions  in  a  solution. 

Mineral  Waters. — Rivers  flowing  into  the  ocean  carry 
small  amounts  of  salt  and  other  soluble  compounds.  Dur- 
ing millions  of  years  water  has  evaporated  from  the  surface 
of  the  ocean  and  the  vapor  has  been  carried  back  over  the 
land  where  it  condenses  and  falls  as  rain,  carrying  a  new 
portion  of  mineral  matters  to  the  rivers  and  the  ocean. 
Some  of  the  substances  carried  to  the  ocean  are  only  slightly 
soluble  and  these  cannot  accumulate  to  any  large  extent,  but 


68  PROPERTIES  AND  USES  OF  WATER 

salt  and  some  other  compounds  are  so  easily  soluble  that  the 
ocean  wa.ter  is  very  far  from  a  saturated  solution  of  them. 
In  the  Dead  Sea,  the  Caspian  Sea,  Salt  Lake  and  other 
inland  bodies  of  water  with  no  outlet  to  the  ocean  the  accu- 
mulation of  soluble  salts  has  gone  much  further;  so  far,  in- 
deed, that  even  very  soluble  salts  may  be  deposited  from 
them.  During  the  long  periods  of  geologic  time  vast  beds 
of  such  salts  have  been  deposited  and  this  has  given  rise 
to  the  beds  of  rock  salt  found  in  New  York,  Kansas,  South- 
eastern Germany  and  elsewhere.  In  some  cases  potash 
salts  and  other  salts  have  been  deposited  and  this  is  espe- 
cially true  in  Stassfurt,  Prussia,  and  in  Alsace.  For  many 
years  the  supply  of  potassium  compounds  for  the  world 
came  almost  exclusively  from  Germany. 

The  waters  of  Searle's  Lake  in  California  and  of  some  lakes 
in  Nebraska  contain  potassium  salts,  and  during  the  great 
war  the  recovery  of  compounds  of  potassium  from  these 
sources  has  been  rapidly  developed.  The  prices  charged 
for  these  salts  are,  however,  much  higher  than  the  prices 
charged  for  potassium  salts  from  Germany  before  the  war. 

Water  which  comes  in  contact  with  deposits  of  salt  will, 
of  course,  dissolve  the  soluble  salts  and  in  this  way  the 
brines  from  which  salt  is  made  are  formed  at  Syracuse, 
New  York;  Midland,  Michigan;  and  elsewhere. 

Water  containing  other  salts  in  solution  may  have  valu- 
able medicinal  properties  and  such  waters  are  sold  as  mineral 
waters.  Artificial  waters  similar  in  composition  are  often 
made  and  sold.  Some  of  these,  as  Appolinaris  water,  are 
charged  with  carbon  dioxide,  to  give  them  their  effervescent 
qualities. 

Hydrates.  Efflorescence,  Deliquescence. — Many  salts 
which  are  separated  from  solution  in  crystalline  form  by 
evaporation  combine  with  definite  quantities  of  water. 
Crystallized  salts  of  this  kind  are  called  hydrates.  Thus 
the  hydrate  of  copper  sulfate  has  the  composition  CuSO4.- 


EFFLORESCENCE.     DELIQUESCENCE  69 

5H20.     The  hydrate  of  sodium  sulfate  is  Na2S04.10H2O. 
One  of  the  hydrates  of  calcium  chloride  is  CaCl2.6H2O. 

Sodium  sulfate,  Na2S04.10H20,  holds  its  water  very 
loosely  and  if  exposed  to  ordinary  air,  which  is  not  saturated 
with  moisture,  the  water  of  the  salt  will  escape  and  the 
partially  dried  salt  will  fall  to  a  powder.  This  is  called 
efflorescence.  The  salts  are  said  to  be  efflorescent.  Other 
salts,  such  as  either  anhydrous  calcium  chloride,  CaCl2,  or 
the  hydrate,  CaCl2.6H2O,  not  only  do  not  lose  water  in 
ordinary  air  but  they  will  even  take  up  more  water  and 
finally  dissolve  in  the  water  absorbed,  unless  the  air  is  very 
dry  indeed.  Salts  of  this  type  are  called  deliquescent. 

SUMMARY 

The  latent  heat  of  fusion  of  ice  is  79.63  calories. 

The  latent  heat  of  evaporation  of  water  is  536.6  calories. 

Water  has  its  maximum  density  at  4°. 

When  water  and  its  vapor  are  in  contact  the  vapor  has  a 
definite  pressure  which  depends  only  on  the  temperature. 
The  temperature  at  which  this  vapor  pressure  is  equal  to 
the  pressure  of  the  air  is  the  boiling  point. 

Equilibrium  between  a  liquid  and  its  vapor  is  reached 
when  the  number  of  molecules  escaping  from  the  liquid  to 
the  vapor  in  a  given  time  equals  the  number  of  molecules 
returning  to  the  liquid  from  the  vapor. 

Germs  of  such  diseases  as  typhoid  fever  and  cholera  may 
be  carried  by  water.  These  may  be  removed  by  proper 
filtration  or  killed  by  boiling  the  water  or  by  sterilization. 

A  saturated  solution  contains  all  of  a  substance  which  the 
liquid  can  retain  while  in  contact  with  the  solid  phase. 
Supersaturated  solutions  may  be  prepared  in  the  absence  of 
the  solid  phase. 

Substances  in  solution  often  react  more  readily  than 
others,  partly  because  they  are' in  intimate  contact,  partly 


70  PROPERTIES  AND  USES  OF  WATER 

because  many  substances  separate  into  ions  in  solutions, 
and  the  ions  unite  to  form  different  compounds. 

In  electrolysis  the  positive  ions,  cations,  move  toward  the 
cathode,  the  negative  ions,  anions,  toward  the  anode.  The 
rate  of  migration  is  different  for  different  ions. 

Mineral  waters  are  found  in  many  places  and  are  some- 
times valuable  for  their  medicinal  properties. 

Many  salts  and  sometimes  other  compounds  combine 
with  water  to  form  hydrates.  Some  of  these  are  efflo- 
rescent, others  deliquescent. 

EXERCISES 

1.  What  per  cent  of  water  is  contained  in  crystallized  copper 
sulfate?     In  crystallized  sodium  sulfate?     In  crystallized  calcium 
chloride? 

2.  How  many  grams  of  copper  will  be  required  to  give  a  pound 
of  crystallized   copper  sulfate?     Of  anhydrous   copper  sulfate? 

3.  What  method  is  used  in  your  community  to  secure  a  water 
supply  free  from  disease  germs? 

4.  The  following  methods  of  purifying  or  sterilizing  water  have 
been  used: 

(a)  Slow  sand  nitration. 

(6)  Addition  of  aluminium  sulfate  (so-called nitration  "alum  ") 
followed  by  rapid,  mechanical  nitration. 

(c)  Aluminium    sulfate    and    "lime"     (calcium   hydroxide) 

followed  by  settling  basins  and  filtration. 

(d)  Ozone. 

(e)  Ultra-violet  light. 
(/)  Bleaching  powder. 
(g)    Liquid  chlorine. 

What  are  some  of  the  advantages  and  disadvantages  of  each  of 
these  methods? 

5.  The  amount  of  aluminium  sulfate  used  in  filtration  plants  is 
from  0.3  to  2  grains  per  gallon.     How  many  gallons  can  be  puri- 
fied by  the  use  of  a  pound?     One  pound  is  7000  grains  and  one 
gallon  is  8H    (8.3389)   pounds. 


EXERCISES.     WATER  71 

6.  From  0.2  to  0.5  parts  per  million  of  liquid  chlorine  are  used 
to  sterilize  water.     How  many  gallons  can  be  sterilized  by  a  pound? 
How  many  gallons  can  be  sterilized  by  a  pound  of  bleaching  powder 
containing  33^  per  cent  of  available  chlorine? 

7.  What  has  been  the  effect  of  the  drainage  canal  on  the  typhoid 
death  rate  in  Chicago? 

8.  How  many  deaths  were  there,  per  thousand  soldiers,  in  the 
cantonments   of   the    Spanish-American    War?     In    the    Russo- 
Japanese  War?     In  the  Great  War?     What  are  the  causes  of  these 
differences? 

9.  By  what  means  other  than  contaminated  water  is  typhoid 
fever  communicated?     What  means  were  used  to  prevent  typhoid 
fever  in  the  armies  of  the  Great  War? 


CHAPTER  VII 
gODIUM,  ACIDS,  BASES,  SALTS 

Salt. — Common  salt  is  a  substance  composed  of  sodium, 
Na,  and  chlorine,  Cl,  and  has  the  composition  represented 
by  the  formula  NaCl.  The  chemical  name  is  sodium 
chloride.  Salt  is  very  widely  diffused.  It  is  a  necessary 
constituent  of  the  food  which  we  eat  and  is  found  in  small 
amounts  in  practically  all  natural  waters  which  have  come 
In  contact  with  the  earth.  The  salty  taste  of  sea  water  is 
cniefly  due  to  sodium  chloride,  though  sea  water  also  con- 
tains a  number  of  other  chlorides.  Thick  beds  of  rock  salt 
are  found  in  Germany,  Louisiana,  Kansas  and  many  other 
places.  Strong  brines,  which  are  nearly  saturated  solutions 
of  sodium  chloride,  are  also  obtained  by  means  of  artesian 
wells, 'and  salt  is  manufactured  by  evaporating  these  brines 
at  Syracuse,  New  York,  in  Michigan,  and  at  many  other 
places  in  the  world. 

Electrolysis  of  Fused  Sodium  Chloride. — If  salt  is  melted 
in  a  crucible  and  an  electric  current  is  passed  through  it 
between  carbon  electrodes,  chlorine  gas  will  be  evolved  at 
the  anode  ancj  sodium  will  be  liberated  at  the  cathode. 

If  a  current  is  passed  through  a  solution  of  sodium 
chloride  in  water,  the  chloride  ions,  Cl~~,  will  be  carried 
toward  the  anoole  by  the  current  and  the  sodium  ions,  Na+, 
will  be  carried  toward  the  cathode.  In  contact  with  the 
anode  the  chloride  ions  will  be  discharged  and  chlorine 
gas  will  be  evolved.  In  any  aqueous  solution  there  are 
present,  besides  the  ions  of  the  dissolved  salt,  hydrogen 
ions  H+,  and  hydroxide  ions,  OH~,  which  come  from  the 
/(--U  72 


SODIUM.     ACIDS  73 

ionization  of  the  water,  H2O  (or  HOH).  Hydrogen  ions 
require  much  less  electrical  energy  than  sodium  ions  for 
their  discharge.  At  the  cathode,  therefore,  the  hydrogen 
ions  are  discharged  and  hydrogen  gas  is  liberated,  while 
the  hydroxide  ions,  OH~,  and  sodium  ions,  Na+,  remain  in 
solution,  forming  a  solution  of  sodium  hydroxide,  NaOH. 
This  method  is  now  extensively  used  for  the  manufacture 
of  sodium  hydroxide  and  also  of  chlorine. 

Sodium  is  a  soft  metal  which  looks  very  much  like  silver 
when  it  is  first  cut.  The  surface  tarnishes  almost  instantly 
in  the  air,  however,  because  it  decomposes  the  moisture  in 
the  air,  liberating  hydrogen  and  forming  sodium  hydroxide : 

2Na  +  2HOH  =  2NaOH  +  H2 

The  same  reaction  occurs  when  a  small  piece  of  sodium 
is  thrown  on  water.  The  metal  is  lighter  than  water  and 
floats  on  the  surface,  melting  to  a  globule  from  the  heat  of 
the  reaction.  This  globule  glides  rapidly  over  the  surface 
of  the  water  as  hydrogen  is  evolved  at  the  surface  of  contact. 
The  sodium  hydroxide,  NaOH,  dissolves  in  the  water  and 
gives  to  it  an  alkaline  reaction,  that  is,  it  will  turn  red 
litmus  blue,  and  it  will  neutralize  acids  (p.  75). 

Acids. — In  a  previous  chapter  several  acids,  such  as 
hydrochloric  acid,  HC1,  sulfuric  acid,  H2SO4,  and  acetic 
acid,  HC2H3O2,  were  spoken  of.  It  was  pointed  out  that 
the  element  common  to  all  acids  is  hydrogen  and  that  this 
hydrogen  may  be  replaced  by  metals,  as  in  the  action  of 
sulfuric  acid  on  zinc,  which  gives  zinc  sulfate,  ZnSO4,  and 
hydrogen. 

Another  very  important  characteristic  of  acids  is  that 
they  separate,  or  ionize,  in  solutions  in  water,  giving  hydro- 
gen ions  which  have  a  positive  charge  and  some  negative 
atom  or  group  of  atoms.  Thus  hydrochloric  acid  gives 
hydrogen  ions,  H+,  and  chloride  ions,  Cl~;  sulfuric  acid 
gives  hydrogen  ions,  H+,  H+,  and  sulfate  ions,  S04=;  acetic 


74  SODIUM,  ACIDS,  BASES,  SALTS 

acid,  HC2H3O2,  gives  hydrogen  ions,  H+,  and  acetate  ions, 
C2H3O2~.  If  a  current  of  electricity  is  passed  through  a 
solution  of  an  acid,  the  hydrogen  ions  are  carried  by  the 
current  toward  the  cathode,  or  negative  electrode,  while 
the  negative  ions  are  carried  toward  the  anode,  or  positive 
electrode. 

Bases. — Sodium  hydroxide,  NaOH,  belongs  to  a  class  of 
compounds  called  bases.  These  ionize  in  solution,  giving 
hydroxide  ions,  OH~,  carrying  a  negative  charge.  Thus 
sodium  hydroxide  gives  hydroxide  ions,  OH~,  and  sodium 
ions,  Na+;  calcium  hydroxide,  Ca(OH)2,  gives  hydroxide 
ions,  OH~,  OH~,  and  calcium  ions,  Ca++,  ammonium  hy- 
droxide, NH4OH,  gives  hydroxide  ions,  OH~,  and  ammonium 
ions,  NH4+. 

From  the  point  of  view  just  presented  the  most  essential 
characteristic  of  an  acid  is  that  it  gives  hydrogen  ions  in 
solution,  and  the  most  characteristic  property  of  a  base  is 
that  its  solution  contains  hydroxide  ions. 

lonization  of  Water.  Neutralization.  Salts. — Pure  water, 
H20,  or  HOH,  separates  to  a  very  slight  extent  into  hydro- 
gen ions,  H+,  and  hydroxide  ions,  OH~.  This  may  be  rep- 
resented by  the  equilibrium  equation: 

HOH  *=>  H+  +  OH- 

In  this  reaction  the  equilibrium  is  very  far  to  the  left, 
i.e.,  there  are  always  a  very  large  number  of  molecules  of 
water  in  comparison  with  the  number  of  hydrogen  and 
hydroxide  ions  (in  pure  water  approximately  500,000,- 
000  :  1)  If  a  solution  of  an  acid,  containing  hydrogen 
ions,  H+,  is  added  to  a  solution  of  a  base,  containing  hydrox- 
ide ions,  OH~,  the  two  kinds  of  ions  must  unite  with  each 
other  until  either  the  hydrogen  or  the  hydroxide  ions,  or 
both,  nearly  disappear.  If  the  acid  and  base  are  used  in 
equivalent  amounts,  the  resulting  solution  must  contain 
equal  numbers  of  hydrogen  and  hydroxide  ions  and  the 


NEUTRALIZATION  75 

number  of  each  of  these  ions  must  be  very  small.  Such 
a  solution  is  said  to  be  neutral.  A  base  is  said  to  neutralize 
an  acid  because  each  destroys  the  characteristic  property 
of  the  other  by  the  process  just  described.  The  process  of 
neutralization  gives  rise  to  reactions  which  may  be  illus- 
trated by  such  equations  as  the  following: 

NaOH  +  HC1  =  NaCl  +  H2O 

Sodium 
chloride 

2NaOH  +  H2SO4  =  Na2S04  +  2H20 

Sodium 
sulfate 

Ca(OH)2  +  2HC2H302  =  Ca(C2H3O2)2  +  2H2O 

Calcium 
acetate 

2A1(OH)3  +  3H2SO4  =  A12(SO4)3  +  3H2O 

Aluminium 
sulfate 

It  is  to  be  noticed  that  in  each  equation  the  number  of 
hydrogen  atoms  in  the  acid  must  be  the  same  as  the  number 
of  hydroxide  groups  in  the  base  and  that  this  fixes  the  for- 
mula of  the  salt. 

The  compound  formed  by  the  union  of  the  positive  ion 
of  a  base  and  the  negative  ion  of  an  acid  is  called  a  salt. 
A  salt  is  also  frequently  defined  as  a  compound  formed  by 
the  replacement  of  the  hydrogen  of  an  acid  by  a  metal. 

Exercise. — Write  the  sixteen  equations  representing  the 
reactions  which  may  occur  between  the  following  acids 
and  bases: 

Hydrochloric  acid,  HC1  Sodium  hydroxide,  NaOH 

Nitric  acid,  HN03  Ammonium  hydroxide,  NH4OH 

Sulfuric  acid,  H2S04  Calcium   hydroxide,  Ca(OH)2 

Phosphoric  acid,  H3PO4  Ferric  hydroxide,  Fe(OH)3 

Indicators. — A  number  of  natural  and  artificial  dyes 
exhibit  a  different  color  in  an  acid  from  that  shown 'in  an 


76  SODIUM,  ACIDS,  BASES,  SALTS 

alkaline  solution.  Thus  in  an  acid  solution,  that  is,  in  a 
solution  which  contains  more  than  a  very  few  hydrogen 
ions,  litmus  is  red,  while  in  an  alkaline  solution  it  becomes 
blue.  Phenol  phthalein,  on  the  other  hand,  is  red  in  an 
alkaline  solution  but  colorless  in  an  acid  solution.  The 
change  in  color  is  due  to  a  change  in  composition  of  the  dye, 
but  the  nature  of  the  change  cannot  be  given  here.  Dyes 
of  this  character  are  called  indicators,  and  they  are  much 
used  to  determine  the  " end-point'*  when  an  acid  is  neu- 
tralized by  a  base.  Paper  which  has  been  dipped  in  a  so- 
lution of  litmus  and  dried  is  used  to  determine  whether  a 
given  solution  is  acid  or  alkaline. 

Dibasic  Acids. — Either  one  or  both  of  the  hydrogen  atoms 
of  sulfuric  acid  may  be  replaced  by  a  metal,  giving  acid  and 
normal  salts,  as  acid  sodium  sulfate,  NaHS04,  and  normal 
sodium  sulfate,  Na2SO4.  Acids  having  this  property  are 
called  dibasic.  An  acid  like  phosphoric  acid,  H3PO4, 
which  forms  three  salts  with  sodium,  NaH2PO4,  Na2HP04 
and  Na3P04,  is  called  tribasic.  The  basicity  depends, 
however,  not  on  the  number  of  hydrogen  atoms  in  one 
molecule  of  an  acid,  but  on  the  number  of  replaceable 
hydrogen  atoms.  Thus  acetic  acid,  C2H4O2,  is  monobasic 
because  only  one  of  its  hydrogen  atoms  can  be  replaced; 
and  phosphorous  acid,  H3PO3,  is  dibasic  because  only  two 
of  the  hydrogen  atoms  can  be  replaced. 

Normal,  Neutral  and  Acid  Salts. — When  a  solution  of  a 
base  is  mixed  with  an  equivalent  amount  of  a  solution  of 
an  acid  the  hydrogen,  H+,  and  hydroxide,  OH~~,  ions  unite  to 
form  water,  H20.  The  metallic  ions  of  the  base,  such  as 
Na+,  from  NaOH,  and  the  non-metallic  ions  of  the  acid, 
such  as  Cl~,  from  hydrochloric  acid,  HC1,  or  SO4=  from  sul- 
furic acid,  H2SO4,  partly  remain  in  solution  as  ions  and 
partly  unite  to  form  salts.  In  the  cases  referred  to  the 
salts  are  sodium  chloride,  NaCl,  and  sodium  sulfate,  Na2S04. 
If  the  solution  is  evaporated  to  dryness  the  ions  will  all 


SALTS  OF  WEAK  ACIDS  77 

unite  to  form  the  salt.  In  other  words,  in  the  reversible 
reactions  represented  by  the  equations: 

NaCl  ±=>  Na+  +  Gl- 
and: 

Na2SO4  *=>  Na+  +  Na+  +  SO4= 

the  separation  into  ions  is  favored  by  the  addition  of  more 
and  more  water,  and  electrolytes,  such  as  the  salts  named, 
and  strong  acids,  like  hydrochloric  or  sulfuric  acid,  are 
almost  completely  ionized  in  dilute  solutions.  On  the  other 
hand,  the  concentration  of  the  solution  by  the  removal  of 
water  causes  the  ionization  to  decrease  and  in  the  dry  salt 
very  few,  if  any,  ions  remain. 

Dibasic  acids  form  acid  salts  such  as  acid  sodium  sulfate, 
NaHSO4,  and  normal  salts,  as  sodium  sulfate,  Na2S04. 
Normal  salts  of  strong  acids,  that  is,  of  acids  which  are 
largely  ionized  in  dilute  solution,  with  strong  bases,  are 
neutral  in  their  reaction  toward  litmus  and  other  indicators. 
This  means,  of  course,  that  the  numbers  of  hydrogen, 
H+,  and  hydroxide,  OH  —  ,  ions  in  solutions  of  such  salts 
are  equal. 

Normal  salts  of  weak  acids,  such  as  carbonic  acid,  H2CO3, 
on  the  other  hand,  are  frequently  'alkaline  toward  litmus. 
Sodium  carbonate,  Na2CO3,  ionizes  in  solution  to  sodium 
ions,  Na+,  Na+,  and  carbonate  ions,  C03=.  The  carbonate 
ions,  however,  combine  with  the  hydrogen  ions  of  the  water 
to  form  bicarbonate  ions,  HC03~: 

CO3=  +  H+  <=>  HCO-r 

Since  the  bicarbonate  ions  have  only  a  slight  tendency  to 
separate  into  hydrogen  ions  and  carbonate,  CO3=,  ions,  be- 
cause carbonic  acid  is  a  very  weak  acid,  this  reaction  removes 
a  considerable  number  of  hydrogen  ions  from  the  solution. 
These  will  be  replaced  by  the  further  ionization  of  water: 

HOH<=±H++  OH- 


78  SODIUM,  ACIDS,  BASES,  SALTS 

The  process  leaves  an  excess  of  hydroxide  ions,  OH~,  in 
the  solution  and  the  reaction  is,  therefore,  alkaline. 

Because  of  these  properties  of  salts  of  weak  acids  those 
salts  in  which  all  of  the  replaceable  hydrogen  of  an  acid  has 
been  displaced  by  a  metal  are  called  normal  salts,  rather 
than  neutral  salts. 

Some  confusion  is  liable  to  arise  because  salts  in  which  a 
part  only  of  the  hydrogen  of  an  acid  has  been  replaced  are 
called  add  salts,  although  such  salts  may  be  neutral  or  even 
alkaline  in  reaction.  Thus  NaHCO3  is  called,  frequently, 
acid  sodium  carbonate,  though  it  is  neutral  in  reaction  to 
litmus.  This  use  of  the  term  add  salt  is  based  on  the  older 
definition  of  an  acid  as  a  compound  containing  replaceable 
hydrogen,  and  is  not  consistent  with  the  modern  definition 
of  an  acid  as  a  compound  whose  solution  contains  more  hydro- 
gen than  hydroxide  ions. 

SUMMARY 

Common  salt  or  sodium  chloride  is  our  most  common 
easily  soluble  salt. 

The  electrolysis  of  melted  salt,  NaCl,  gives  metallic 
sodium  and  chlorine. 

The  electrolysis  of  a  solution  of  salt  gives  sodium  hydrox- 
ide and  hydrogen  at  the  cathode  and  chlorine  at  the  anode. 

Acids  are  hydrogen  compounds  which  ionize  to  positive 
hydrogen  and  negative  ions.  They  are  also  defined  as 
hydrogen  compounds  in  which  the  hydrogen  may  be  re- 
placed by  a  metal. 

Bases  are  hydroxyl  (OH)  compounds  which  ionize  to 
negative  hydroxide  ions  and  some  positive  ion,  which  is 
usually  metallic. 

Neutralization  consists  in  the  union  of  the  hydrogen  ions 
of  an  acid  with  the  hydroxide  ions  of  a  base,  forming  water, 
while  the  negative  ion  of  the  acid  and  the  positive  ion  of  the 


EXERCISES.     ACIDS,   BASES  79 

base  either  remain  in  solution  uncombined  or  unite  to  form  a 
salt. 

In  a  neutral  solution  the  numbers  of  hydrogen  and  hydrox- 
ide ions  are  equal. 

An  indicator  exhibits  one  color  in  the  presence  of  an  ex- 
cess of  hydrogen  ions  and  another  color  in  the  presence  of 
an  excess  of  hydroxide  ions. 

Normal  salts  are  salts  in  which  all  of  the  replaceable 
hydrogen  of  an  acid  has  been  replaced  by  a  metal  or  some 
other  metallic  group.  Normal  salts  may  be  neutral,  alkaline 
or  acid  in  reaction. 

Acid  salts,  in  the  older  use  of  the  term,  are  salts  in  which 
only  a  part  of  the  replaceable  hydrogen  has  been  replaced 
by  a  metal. 

EXERCISES 

1.  How  many  grams  of  hydrochloric  acid  must  there  be  in  a 
liter  of  a  solution  of  the  acid  which  will  exactly  neutralize  a  liter 
of  a  solution  containing  40  grams  of  sodium  hydroxide? 

2.  How  many  grams  of  sulfuric  acid  in  one  liter  will  give  a 
solution   exactly  equivalent  to   the  sodium  hydroxide  solution 
just  mentioned?     How  many  grams  of  nitric  acid?     Of  acetic 
acid? 

3.  How  much  lime,  CaO,  will  be  neutralized  by  one  liter  of  any 
one  of  the  acid  solutions  mentioned  above? 

4.  What-is  the  difference  between  the  method  of  preparing  salts 
given  in  this  chapter  and  that  given  in  Chapter  IV? 

5.  Write  the  equation  for  the  reaction  between  acetic  acid  and 
calcium.     For  the  reaction  between  sulfuric  acid  and  sodium. 

6.  Why  does  a  globule  of  melted  sodium  move  rapidly  over  the 
surface  of  water? 


CHAPTER  VIII 

HYDROCHLORIC  ACID.     CHLORINE.     OXYGEN 
ACIDS  OF  CHLORINE 

Preparation  of  Hydrochloric  Acid. — The  addition  of  con- 
centrated sulfuric  acid,  H2S04,  to  salt,  NaCl,  causes  the 
liberation  of  hydrochloric  acid  as  a  gas : 

NaCl  +  H2S04  ?=*  NaHSO4  +  HG1 

Sodium  Sulfuric  Acid  Hydrochloric 

chloride  acid  sodium  acid 

sulfate 

The  compound  NaHSO4  is  called  acid  sodium  sulfate  be- 
cause it  still  has  acid  properties.  How  might  this  be  shown? 

It  is  sometimes  stated  that  hydrochloric  acid  is  liberated 
because  sulfuric  acid  is  a  stronger  acid  than  hydrochloric- 
Such  a  statement  overlooks  the  fact  that  the  addition  of 
concentrated  hydrochloric  acid  to  a  concentrated  solution 
of  acid  sodium  sulfate  causes  a  precipitation  of  salt,  NaCl. 

NaHSO4  +  HC1  <=>  NaCl  +  H2SO4 

If  sulfuric  acid  is  the  stronger  acid  in  one  case,  hydro- 
chloric acid  is  stronger  in  the  other.  It  is  clear,  therefore, 
that  the  direction  of  the  reaction  does  not  depend  on  the 
relative  strengths  of  the  two  acids.  It  depends,  instead, 
on  the  fact  that  the  reaction  is  reversible  and  in  the  first 
case  it  goes  toward  the  liberation  of  hydrochloric  acid 
because  that  escapes  as  a  gas,  and  in  the  second  case  it  goes 
toward  the  formation  of  salt  because  the  salt  is  removed 
from  the  sphere  of  action  as  a  precipitate. 

Properties  of  Hydrochloric  Acid. — Hydrochloric  acid  is  a 
colorless  gas,  about  one-fourth  heavier  than  air.  It  has  a 

80 


ACTION  ON  METALS  81 

very  pungent,  irritating  odor,  even  when  diluted  with  a 
large  amount  of  air.  It  dissolves  in  water  and  at  the  freez- 
ing point  one  volume  of  water  will  take  up  500  volumes 
of  the  gas.  The  solution  contains  forty-five  per  cent  of 
hydrochloric  acid. 

This  concentrated  solution  gives  off  some  of  the  gas  at 
ordinary  temperatures.  If  the  solution  is  boiled,  hydro- 
chloric acid  escapes  chiefly  at  first  and  the  solution  becomes 
less  and  less  concentrated,  till  there  finally  remains  a 
solution  which  contains  only  twenty  per  cent  of  hydro- 
chloric acid  and  which  boils  at  about  110°.  On  the  other 
hand,  a  dilute  solution  containing  less  than  twenty  per 
cent  of  the  acid  allows  water  to  escape  at  first,  when 
boiled,  and  the  boiling  point  gradually  rises  to  110°,  and 
the  concentration  of  the  acid  remaining  increases  to  twenty 
per  cent.  This  acid,  which  boils  constantly  at  110°,  may 
be  obtained,  as  described,  from  either  a  more  concentrated 
or  from  a  more  dilute  acid. 

When  hydrochloric  acid  escapes  into  the  air  a  cloud  is 
formed  unless  the  air  is  very  dry.  This  is  because  the  acid 
forms  with  the  water  of  the  air  a  mixture  or  compound 
having  a  higher  boiling  point  than  the  boiling  point  of 
pure  water.  This  mixture  will  condense  to  minute  drops 
of  a  concentrated  solution  of  hydrochloric  acid  at  a  tem- 
perature at  which  the  water  contained  in  the  volume  of  air 
occupied  by  the  cloud  would  all  remain  as  an  invisible 
vapor. 

Action  of  Hydrochloric  Acid  on  Metals. — The  solution 
of  hydrochloric  acid  in  water  acts  on  many  of  the  metals, 
forming  chlorides  of  the  metals  while  the  hydrogen  of  the 
acid  is  liberated  as  a  gas.  In  this  way  sodium  gives  sodium 
chloride,  NaCl;  zinc  gives  zinc  chloride,  ZnCl2;  iron  gives 
ferrous  chloride,  FeCl2;  aluminium  gives  aluminium 
chloride,  A1C13,  and  tin  gives  stannous  chloride,  SnCl2. 


82  HYDROCHLORIC  ACID 

The  equations  for  these  reactions  should  be  written  by 
the  student. 

Action  of  Hydrochloric  Acid  on  Hydroxides  and  Oxides 
of  the  Metals. — The  formation  of  salts  by  the  Action  of 
acids  on  bases  was  discussed  in  the  last  chapter  and  should 
be  recalled  here. 

In  a  somewhat  similar  manner  many  oxides  of  the  metals 
dissolve  in  a  solution  of  hydrochloric  acid,  giving  chlorides 
of  the  metals  and  water.  Thus  zinc  oxide  gives  zinc 
chloride : 

ZnO  +  2HC1  =  ZnCl2  +  H2O 

Zinc  Zinc 

oxide  chloride 

It  is  to  be  noticed  that  in  reactions  of  this  character  two 
molecules  of  hydrochloric  acid  are  required  for  each  atom 
of  oxygen  in  the  oxide.  In  a  similar  manner  ferric  oxide, 
Fe2O3,  cuprous  oxide,  Cu2O,  cupric  oxide,  CuO,  mercuric 
oxide,  HgO,  and  bismuth  oxide,  Bi2O3,  give  corresponding 
chlorides.  The  magnetic  oxide  of  iron,  Fe3O4,  gives  a  mix- 
ture of  ferrous  chloride,  FeCl2,  and  ferric  chloride,  FeCl3. 
The  student  should  write  the  equations  for  all  of  these 
reactions. 

Action  of  Oxidizing  Agents  on  Hydrochloric  Acid. — By 
heating  a  mixture  of  hydrochloric  acid  with  air  or  oxygen  a 
part  of  the  hydrogen  of  the  acid  is  oxidized  to  water  and 
chlorine  is  liberated.  The  reaction  is  reversible  and  may  be 
represented  thus: 

4HC1  +  O2  ±=F  2H2O  +  2C12 

By  using  copper  chloride  as  a  catalytic  agent  (p.  9)  the 
reaction  has  sometimes  been  used  for  the  commercial 
preparation  of  chlorine,  but  as  a  technical  process  this 
has  not  been  able  to  compete  with  other  methods  of  obtain- 
ing the  element. 

The  oxidation  of  hydrochloric  acid  may  also  be  effected 


PROPERTIES  OF  CHLORINE  83 

by  a  number  of  oxidizing  agents  at  ordinary  temperatures, 
or  by  gently  warming  a  concentrated  solution  of  hydrochloric 
acid  with  the  agent.  The  substances  most  often  used  in 
the  laboratory  for  this  purpose  are  manganese  dioxide, 
MnO2,  potassium  permanganate,  KMnO4,  and  potassium 
dichromate,  K2Cr2O7.  In  the  first  case  manganese  chloride, 
MnCl2,  is  formed,  in  the  second  case  manganese  chloride 
and  potassium  chloride: 

Mn02        +  4HC1    =  MnCl2    +  2H2O  +  C12 
2KMnO4  +  16HC1  =  2MnCl2  +  2KC1  -f  8H2O  +  5C12 

It  should  be  noticed  that  in  each  case  enough  molecules 
of  hydrochloric  acid  must  be  used  to  furnish  hydrogen  to 
combine  with  all  of  the  oxygen  of  the  oxidizing  agent. 

Properties  of  Chlorine. — Chlorine  forms  compounds  with 
nearly  all  of  the  chemical  elements,  and  in  many  cases  the 
action  of  chlorine  upon  an  element  takes  place  at  a  much 
lower  temperature  than  that  at  which  the  same  element 
will  combine  with  oxygen.  Finely  powdered  antimony 
will  burst  into  flame  when  sprinkled  through  the  gas,  and 
false  gold  leaf,  an  alloy  of  copper  and  zinc,  will  also  flash 
up  when  the  gas  is  passed  over  it.  The  antimony  forms 
antimony  trichloride,  SbCl3,  or  antimony  pentachloride, 
SbCl5,  according  as  the  antimony  or  chlorine  is  in  excess. 
Phosphorus  forms  similar  compounds,  phosphorus  tri- 
chloride, PC13,  and  phosphorus  pentachloride,  PC15.  The 
copper  and  zinc  of  the  false  gold  leaf  form  cuprous  chloride, 
Cu2Cl2,  or  cupric  chloride,  CuCl2,  and  zinc  chloride,  ZnCl2. 

Chlorine  gas  is  very  poisonous  and  great  care  should  be 
taken  to  avoid  breathing  it.  When  it  has  been  inhaled 
in  small  quantities  the  best  antidote  for  immediate  use 
is  vapor  of  alcohol.  The  gas  masks  used  in  warfare  contain 
soda-lime  and  sodium  thiosulfate.  For  the  absorption 
of  other  poisonous  gases  charcoal  prepared  from  cocoanut 
shells  is  used. 


84  HYDROCHLORIC  ACID 

Bleaching. — Many  vegetable  and  artificial  colored  sub- 
stances are  bleached  by  moist  chlorine  gas  or  by  a  solution 
of  chlorine  in  water.  The  brown  colors  which  are  charac- 
teristic of  unbleached  cotton  and  linen  are  also  destroyed 
in  this  way.  These  colors  are  sensitive  to  sunlight,  and  be- 
fore the  method  of  bleaching  with  chlorine  was  discovered 
it  was  customary  to  spread  cotton  or  linen  cloth  on  the 
grass  and  expose  it  to  the  action  of  the  sunlight  for  many 
days  or  weeks  to  secure  the  pure  white  color  which  was 
desired.  The  same  effect  is  now  obtained  very  quickly 
by  the  use  of  a  dilute  solution  of  chlorine.  Dry  chlorine 
acts  very  slowly  or  not  at  all  on  the  colored  calico.  Because 
of  this  fact  it  is  believed  that  the  chlorine  does  not  act 
directly  on  the  moist  colored  substances  but  that  it  acts 
at  first  on  the  water,  forming  hypochlorous  acid,  HC10, 
and  hydrochloric  acid: 

C12  +  HOH  =  HC1  +  HOC1 

Hydrochloric     Hypochlorous 
acid  acid 

The  hypochlorous  acid  gives  its  oxygen  readily  to  other 
substances  and  seems  to  be  the  actual  bleaching  agent. 
It  oxidizes  the  colored  compounds  to  other  compounds 
which  are  colorless. 

Names  of  Oxygen  Acids  of  Chlorine  and  of  Their  Salts.— 
Chlorine  combines  with  hydrogen  and  oxygen  to  form  four 
different  acids.  Each  of  these  contains  one  atom  of  hydro- 
gen and  one  atom  of  chlorine  in  its  molecule  and,  with  these, 
one,  two,  three  or  four  atoms  of  oxygen.  The  names  of 
these  acids  are: 

Hypochlorous  acid,  HC10 

Chlorous  acid,  HC102 

Chloric  acid,  HC103 

Perchloric  acid,  HC1O4 

In  these  names  the  endings  ous  and  ic  are  used  as  they 
are  used  in  naming  oxides  or  chlorides  (p.  48),  the  com- 


NOMENCLATURE.     HYPOCHLOROUS  ACID  85 

pound  with  the  ending  ous  containing  less,  and  the  one 
having  the  ending  ic  containing  more  oxygen.  The  pre- 
fix hypo  means  under  or  less,  and  the  prefix  per  means  above 
or  beyond.  Hypochlorous  acid  contains  less  oxygen  than 
chlorous  acid;  perchloric  acid  contains  more  oxygen  than 
chloric  acid. 

Corresponding  to  these  acids  are  many  salts  in  which 
the  hydrogen  of  the  acids  is  replaced  by  metals.  Thus 
the  sodium  salts  are: 

Sodium  hypochlorite,  NaCIO 

Sodium  chlorite,  NaClO2 

Sodium  chlorate,  NaClO3 

Sodium  perchlorate,  NaC104 

In  these  names  the  ous  of  the  acids  is  changed  to  ite  for  the 
salts  and  the  ic  of  the  acids  is  changed  to  ate  for  the  salts. 

In  accordance  with  these  principles  name  the  following 
compounds : 

HBrO,  HBr03,  KBrO,  KBr03,  KIO3,  KIO4,  H2SO3, 

H2S04,  Na2S03,  Na2S04,  H3P03,  H3P04,  K3P04. 

Hypochlorous  Acid.  Effect  of  Removing  One  of  the 
Products  of  a  Reversible  Reaction. — When  chlorine  is  dis- 
solved in  water  a  small  part  of  it  reacts  with  the  water  in 
accordance  with  the  reversible  reaction  expressed  by  the 
equation : 

C1-C1  +  H-OH1  <=>  H-C1  +  HO-C1  (or  HC10) 

Hydrochloric  Hypochlorous 

acid  acid 

In  this  reaction  the  equilibrium,  which  is  always  reached 
sooner  or  later  in  a  reversible  reaction,  is  far  to  the  left.  In 
other  words,  there  will  be  a  much  larger  proportion  of  chlo- 
rine (C12)  in  the  solution  than  there  will  be  of  hydrochloric 

1  The  formulas  of  chlorine,  Cl2,  and  water,  HaO,  are  written  in  this 
manner  to  call  attention  to  the  way  in  which  each  substance  divides  in 
reacting  with  the  other. 


86  HYDROCHLORIC  ACID 

and  hypochlorous  acids,  because  the  reaction  between  the 
hydrochloric  acid  and  hypochlorous  acid  giving  chlorine 
and  water  goes  very  much  faster  than  that  between  chlorine 
and  water  giving  hypochlorous  acid  and  hydrochloric  acid. 
If  a  solution  of  potassium  hydroxide  or  of  any  other 
strong  alkali  is  added  to  such  a  solution,  the  base  reacts 
with  the  acids  in  accordance  with  the  usual  interaction  of 
acids  and  bases: 

HC1     +  KOH  *±  KC1     +  HOH 
HC10  +  KOH  <=»  KC10  +  HOH 

The  reactions  shown  by  these  equations  are  also  reversi- 
ble with  the  equilibrium  very  far  to  the  right,  so  that  very 
little  of  the  acids  can  remain  in  a  solution  containing  a  base. 
As  soon,  however,  as  the  hydrochloric  and  hypochlorous 
acids  are  neutralized  by  the  base  the  chlorine  and  water 
will  react  to  form  more  of  the  acids,  and  all  of  the  chlorine 
will  be  very  quickly  converted  into  a  mixture  of  potassium 
chloride,  KC1,  and  potassium  hypochlorite,  KC1O.  As 
these  substances  are  the  first  tangible  products  of  the 
interaction  of  chlorine  and  potassium  hydroxide  we  may 
add  the  three  equations  together,  omitting  compounds 
which  occur  on  opposite  sides : 

C12  +  HOH  =  HC1  +  HC10 
HC1  +  KOH  =  KC1  +  HOH 
HC10  +  KQH  =  KC1Q  +  HOH 
01,  +  2KOH  =  KC1  +  KC10  +  H2O 

Bleaching  Powder  or  Chloride  of  Lime. — In  calcium 
hydroxide,  Ca(OH)2,  or  slaked  lime,  there  are  two  hydroxide 
groups,  OH,  in  the  same  molecule.  When  chlorine  reacts 

Cl 

with  calcium  hydroxide,  therefore,  a  compound,  Ca<^       , 

CIO 


BLEACHING  POWDER.     CHLORIC  ACID  8? 

is  formed  which  is  partly  chloride  and  partly  hypochlorite: 

OH  Cl 

Ca/        +  C12  =  Ca/         +  H^O 
OH  CIO 

Bleaching 
powder 

Bleaching    powder    is    intermediate    between    calcium 
Cl  CIO 

chloride,    Ca\^     ,   and    calcium    hypochlorite,    Ca<^         : 
Cl  CIO 

Bleaching  powder  or  chloride  of  lime  has  proved  to  be  the 
most  convenient  method  of  transporting  chlorine  for  most 
of  the  technical  uses  to  which  it  is  applied.  It  is  a  solid 
which  can  be  readily  carried  in  bulk  in  ordinary  receptacles 
and  when  chlorine  is  wanted  for  any  purpose  the  addition 
of  an  acid  to  the  bleaching  powder  causes  the  reversal  of 
the  reactions  which  have  been  given  above  and  liberates, 
chlorine : 

Cl 

Ca/         +  H2SO4  =  CaSO4  4  HC1  -f  HC1O 

OC1 

HC1O  -r  HC1  =  C12  +  HOH 

Sulfuric  acid  is  usually  chosen  to  liberate  the  chlorine? 
because  it  is  the  cheapest  of  the  acids. 

Chloric  Acid.  Potassium  Chlorate. — It  will  be  remem-* 
bered  that  potassium  chlorate,  KClOs,  decomposes  easily, 
when  heated,  into  potassium  chloride,  KC1,  and  oxygen. 
This  decomposition  is  used  as  the  simplest  laboratory 
method  of  preparing  oxygen. 

The  hypochlorites  all  decompose,  even  more  easily, 
chlorides  and  oxygen: 

2KC1O  =  2KC1  +  0, 

Potassium  Potassium 

hypochlorite  chloride 


88  HYDROCHLORIC  ACID 

The  hypochlorites  may  also  take  up  oxygen  and  pass  into 
the  more  stable  chlorate: 

KC10  +  2O  =  KC1O3 

Potassium 
chlorate 

In  a  warm  solution,  as,  for  instance,  when  chlorine  gas 
is  passed  into  a  concentrated  solution  of  potassium  hydrox- 
ide, the  two  reactions  may  occur  simultaneously  in  accord- 
ance with  the  equation: 

2KC10  +  KC10  =  KC103  +  2KC1 

Two  molecules  of  potassium  hypochlorite  give  their  oxy- 
gen to  a  third  molecule: 


KC1 
KC1 


O 
0 
KC10 

If  we  add  this  equation  to  three  times  the  equation  given 
above  to  represent  the  action  of  chlorine  on  a  cold  solution 
of  potassium  hydroxide,  canceling  the  potassium  hypo- 
chlorite, KC10,  from  both  sides,  we  have  the  following: 

6KOH  +  3C12  =  3KC10  +  3KC1  +  3H20 

3KC1O  +  =  KC1O3  +  2KC1 

6KOH  +  3C12  =  KClOa  +  5KC1  +  3H20 

This  last  equation  may  be  taken  as  representing  what 
occurs  when  chlorine  gas  is  absorbed  by  a  warm  solution 
of  potassium  hydroxide. 

When  a  strong  acid  is  added  to  a  solution  of  potassium 
chlorate  some  of  the  chloric  acid  is  liberated  in  the  solution : 

2KC103  +  H2SO4  ^  2HC103  +  K2S04 

If  such  a  solution  is  evaporated,  however,  the  chloric  acid 
decomposes  into  water,  chlorine,  chlorine  dioxide,  C1O2, 
and  other  substances,  and  it  has  never  been  found  possible 


POTASSIUM  PERCHLORATE  89 

to  prepare  chloric  acid  as  a  pure  compound,  free  from  water. 
If  concentrated  sulfuric  acid  is  added  to  potassium  chlorate, 
these  decompositions  take  place  at  once,  and  chlorine  dioxide 
may  cause  the  ignition  of  sugar  or  other  organic  matter 
in  contact  with  the  mixture.  Matches  were  once  made 
on  the  basis  of  these  facts  but  they  were  soon  displaced  by 
the  more  convenient  phosphorus  matches. 

Potassium  Perchlorate. — The  decomposition  of  potas- 
sium chlorate  by  heat  for  the  preparation  of  oxygen  will 
be  remembered  (p.  9).  If  potassium  chlorate  is  heated 
just  above  its  melting  point  for  some  time,  one  portion 
of  the  salt  gives  oxygen  to  another  portion,  converting  it 
into  potassium  perchlorate: 

KC103  +  3KC103  =  KC1  +  3KC1O4 

Potassium 
perchlorate 

One  molecule  of  potassium  chlorate  gives  enough  oxygen 
to  convert  three  molecules  of  the  chlorate  to  the  perchlorate. 

Potassium  perchlorate  is  much  more  difficultly  soluble 
in  water  than  potassium  chlorate.  It  requires  a  higher 
temperature  for  its  decomposition  into  potassium  chloride 
and  oxygen.  It  is  noticeable  that  the  acids  of  chlorine 
containing  oxygen  and  their  salts  become  more  and  more 
stable  as  the  amount  of  oxygen  increases.  Potassium 
chlorate  is  very  much  more  stable  than  potassium  hypo- 
chlorite  and  the  perchlorate  is  more  stable  than  the  chlo- 
rate. The  increase  in  stability  is  still  more  marked  in  the 
case  of  the  free  acids. 

Perchloric  Acid,  HC1O4,  may  be  distilled  under  low 
pressures  without  decomposition. 

If  a  solution  of  perchloric  acid  is  added  to  a  solution  of 
potassium  chloride  or  of  some  other  soluble  potassium 
salt,  a  precipitate  will  be  formed  because  potassium  per- 
chlorate is  only  very  slightly  soluble: 

KC1  +  HC1O4  =  KC104  +  HC1 


90  HYDROCHLORIC  ACID 

This  property  of  perchloric  acid  is  used  for  the  detection 
and  also  for  the  quantitative  determination  of  potassium. 

SUMMARY 

Hydrochloric  acid  is  a  gas  which  is  prepared  by  the  action 
of  sulfuric  acid  on  salt. 

Hydrochloric  acid  dissolves  easily  in  water.  A  solution 
containing  20  per  cent  of  the  acid  boils  constantly  at  110°. 
A  very  concentrated  solution  gives  chiefly  hydrochloric  acid 
when  boiled,  while  a  dilute  solution  gives  mainly  water. 

Hydrochloric  acid  gives  with  zinc,  iron,  tin  and  other 
metals  hydrogen  and  a  chloride  of  the  metal. 

Hydrochloric  acid  gives  with  oxides  or  hydroxides  of 
metals  water  and  a  chloride  of  the  metal. 

Oxidizing  agents  oxidize  the  hydrogen  of  hydrochloric 
acid,  liberating  chlorine. 

Chlorine  is  an  active  element  combining  directly  both 
with  metals  and  with  non-metals. 

Moist  chlorine  bleaches  nearly  all  vegetable  colors,  de- 
stroying them. 

A  logical  nomenclature  for  acids  and  salts  is  developed 
by  use  of  the  endings  ous  and  ic,  ite  and  atet  and  the  prefixes 
hypo  and  per. 

The  primary  action  of  chlorine  with  water  gives  hypo- 
chlorous  and  hydrochloric  acids. 

Hypochlorites  are  prepared  by  the  action  of  chlorine  on 
hydroxides,  a  chloride  being  formed  at  the  same  time. 

Chlorine  gives  with  calcium  hydroxide  a  compound  which 
is  both  hypochlorite  and  chloride. 

Potassium  hypochlorite  when  warmed  in  solution  gives 
potassium  chloride  and  potassium  chlorate. 

When  potassium  chlorate  is  heated  to  the  melting  point 
it  is  partly  changed  to  a  mixture  of  potassium  perchlorate 
and  potassium  chloride. 


EXERCISES.     CHLORINE  91 

Potassium  perchlorate  is  precipitated  when  a  solution  of 
perchloric  acid  is  added  to  a  solution  of  a  potassium  salt. 

EXERCISES 

1.  Write  the  equations  for  the  following: 
(a)  Calcium  chloride  and  sulfuric  acid. 

(6)  Chlorine  and  a  cold  solution  of  sodium  hydroxide. 

(c)  Chlorine  and  a  warm  solution  of  sodium  hydroxide.. 

(d)  The  effect  of  boiling  a  solution  of  bleaching  powder. 

(e)  Potassium  sulfate  and  perchloric  acid. 
(/)  Sodium  chlorate  when  heated. 

2.  If  slaked  lime  is  treated  with  chlorine  and  the  water  formed 
does   not   escape,    what  per  cent  of  chlorine  will  the  product 
contain? 

3.  How  many  pounds  of  chlorine  will  be  required  to  make  a  ton 
of  bleaching  powder? 

4.  How  much  manganese  dioxide  will  be  required  to  generate  the 
chlorine? 

5.  How  much  hydrochloric  acid  containing  40  per  cent  of  the 
acid  will  be  required  to  generate  the  chlorine? 

6.  How  much  salt  and  how  much  sulfuric  acid  will  be  required  to 
give  the  hydrochloric  acid? 


CHAPTER  IX 

GROUP  VII:  THE  HALOGEN  FAMILY,  BROMINE, 
IODINE  AND  FLUORINE 

Groups  of  Elements. — Although  the  vast  number  of 
compounds  which  are  known,  with  their  infinite  variety  of 
compositions  and  properties,  are  all  made  by  different 
combinations  of  eighty  or  a  few  more  elements,  even  these 
elements  are  not  substances  which  are  wholly  different  from 
each  other.  There  are  three  elements,  bromine,  iodine  and 
fluorine,  which  closely  resemble  chlorine.  In  a  similar 
manner  there  are  three  elements  with  properties  somewhat 
resembling  those  of  oxygen,  and  the  other  elements  fall  into 
a  series  of  well-defined  groups.  After  studying  carefully 
the  properties  of  one  of  the  elements  of  such  a  group  it  is 
found  that  many  of  these  properties  are  repeated  with  com- 
paratively few  differences  in  the  other  members  of  the 
same  group. 

.  Halogen  Family. — The  group  of  elements  which  includes 
chlorine  is  called  the  halogen  family.  The  word  halogen 
means  salt  former  and  the  name  was  given  because  the  ele- 
ments combine  directly  with  metals  to  form  salts  which 
contain  only  two  elements,  metal  and  halogen.  Common 
salt,  or  sodium  chloride,  is  an  illustration.  The  name  con- 
trasts these  salts  with  the  more  common  salts,  such  as 
sodium  sulfate,  Na2SC>4,  which  contain  oxygen  in  addition 
to  the  metal  and  the  non-metal. 

The  elements  of  the  halogen  family  are : 

Fluorine,       19 
Chlorine,       35.5 
Bromine,       80 
Iodine,         127 

92 


BROMINE  93 

The  atomic  weights  are  given  in  round  numbers.  The 
student  is  advised  to  learn  these  atomic  weights,  as  the 
atomic  weights  of  the  succeeding  groups  of  non-metallic 
elements  are  closely  related  to  these  and  these  numbers  are 
the  best  method  which  has  been  found  for  grouping  the 
elements  and  retaining  a  knowledge  of  the  relations  between 
them  in  mind. 

Bromine. — Sea  water  and  most  of  the  brines  from  which 
salt,  NaCl,  is  obtained  by  evaporation  contain  small 
quantities  of  sodium  bromide,  NaBr,  and  still  smaller 
quantities  of  sodium  iodide,  Nal.  The  electrolysis  of  such 
a  brine  by  passing  an  electric  current  through  it  causes 
the  liberation  of  iodine  first,  then  of  bromine  and  finally 
chlorine.  The  amount  of  iodine  in  most  of  the  brines  is 
very  small  indeed,  but  some  of  the  brines,  in  Michigan, 
especially,  contain  enough  bromine  so  that  the  element  may 
be  extracted  profitably  from  them.  After  passing  an 
electric  current  through  the  brine  till  all  of  the  bromine  and 
a  little  chlorine  are  liberated  the  bromine  is  distilled  away. 
Bromine  is  a  dark  liquid  about  three  times  as  heavy  as 
water.  It  gives  off  a  heavy,  reddish  brown  vapor  which  is 
very  irritating  and  poisonous  and  which  has  a  disagreeable 
odor.  In  fact  the  name  bromine  is  derived  from  a  Greek 
word,  jftpwjuos,  meaning  a  stench. 

Compounds  of  Bromine. — Bromine  forms  many  com- 
pounds which  are  similar  to  the  corresponding  compounds 
of  chlorine.  Among  these  may  be  mentioned  hydrobromic 
acid,  HBr,  sodium  bromide,  NaBr,  potassium  bromide, 
KBr,  hypobromous  acid,  HBrO,  potassium  hypobromite, 
KBrO,  bromic  acid,  HBrO3,  and  potassium  bromate, 
KBrO,. 

Frofti  the  analogy  with  chlorine  answer  the  following 
questions: 

What  will  be  the  effect  of  water  on  hydrobromic  acid 
gas? 


§4  THE  HALOGEN  FAMILY 

Write  the  equations  for  the  reaction  between  hydro- 
bromic  acid  and  manganese  dioxide.  Also  for  that  between 
hydrobromic  acid  and  potassium  permanganate. 

Write  the  equations  for  the  reaction  by  which  potassium 
hypobromite  is  prepared.  Also  the  one  for  the  reaction  by 
Which  potassium  bromate  is  formed. 

What  will  be  formed  when  potassium  bromate  is 
heated? 

Sodium  Bromide,  NaBr,  is  used  in  medicine  to  induce 
sleep.  Potassium  bromide,  KBr,  which  was  used  formerly 
for  the  same  purpose,  is  less  suitable  because  of  the  irritating 
effect  of  the  potassium  which  it  contains. 

Silver  Bromide,  AgBr,  is  used  in  photography  (see  p.  311). 

Iodine.-^-The  occurrence  of  small  amounts  of  iodine  in 
sea  water  and  in  brine?  has  been  mentioned.  It  is  not 
practicable  to  obtain  the  element  directly  from  these 
Sources,  but  some  sea  weeds  accumulate  iodine  in  their 
growth  and  from  the  ash  of  these  weeds,  called  kelp,  the 
iodine  can  be  obtained. 

Free  iodine  crystallizes  in  black,  shining  scales.  It 
hielts  when  heated  gently  and  gives  off  a  beautiful  violet 
Vapor.  The  vapors  condense  as  crystals  on  any  cool 
Surface,  An  evaporation  of  a  solid  without  melting  and 
condensation  of  the  vapor  is  called  sublimation.  Iodine 
melts  below  its  boiling  point,  but  it  also  sublimes  below 
its  melting  point. 

Iodine  dissolves  only  very  slightly  in  water.  The  solution 
in  alcohol  is  called  tincture  of  iodine.  It  is  used  as  an 
application  to  bruises  and  swellings  because  of  its  germicidal 
properties.  The  solution  of  iodine  in  alcohol  is  brown,  but 
that  in  carbon  disulfide  and  in  some  other  solvents  is  violet. 
With  a  solution  of  starch  in  the  presence  of  a  little  hydriodic 
acid  or  an  iodide,  iodine  gives  an  intense  blue  color.  This 
may  be  used  as  a  very  sensitive  test  either  for  iodine  or 
starch. 


IODINE.     FLUORINE  95 

Iodine  forms  many  compounds  similar  to  those  of  bromine 
and  chlorine.  Name  and  give  the  formulas  of  some  of 
those  which  might  be  expected. 

Potassium  Iodide,  KI,  is  used  in  medicine.  An  organic 
compound  containing  iodine  is  found  in  the  thyroid  gland 
of  sheep  and  the  lack  of  this  compound  in  the  human  body 
seems  to  be  a  chief  cause  of  goiter.  This  compound  has 
been  prepared,  recently,  at  the  Mayo  laboratory,  Rochester, 
Minnesota,  and  gives  remarkable  results  when  used  as  a 
medicine. 

Fluorine. — The  compounds  of  chlorine,  bromine  and 
iodine  with  calcium,  calcium  chloride,  CaCl2,  calcium 
bromide,  CaBr2,  and  calcium  iodide,  Cal2,  are  all  easily 
soluble  in  water,  but  the  corresponding  compound  of  fluor- 
ine, CaF2,  is  almost  insoluble.  As  compounds  of  calcium 
are  present  in  practically  all  natural  waters  it  is  impossible 
for  sea  water  or  for  brines  to  contain  more  than  a  very 
small  amount  of  fluorine.  The  element  is  found  chiefly 
in  the  form  of  calcium  fluoride,  CaF2,  as  the  mineral  fluorite 
or  fluorspar.  The  mineral  crystallizes  in  cubes  (or  octa- 
hedra).  The  pure  mineral  is  colorless  but  it  is  often  colored 
violet  or  amethyst  by  the  presence  of  a  very  minute  amount 
of  some  other  compound,  or  substance,  possibly  a  com- 
pound of  manganese. 

Properties  of  Fluorine. — Fluorine  is  the  most  active  of 
the  non-metallic  elements.  While  chlorine  and  other 
elements  of  the  family  can  be  easily  prepared  by  the 
electrolysis  of  solutions  of  their  compounds  or  by  easy 
reactions  which  occur  in  the  presence  of  water,  the  electro- 
lysis of  an  aqueous  solution  of  a  fluoride  gives  oxygen  or 
ozone  in  place  of  fluorine  because  free  fluorine  reacts  with 
water,  liberating  ozone,  Os : 

3F2  +  3H2O  =  6HF  +  O3 

Hydro-    Ozone 
fluorio 
acid 


96  THE  HALOGEN  FAMILY 

It  was  not  till  this  property  of  fluorine  was  clearly 
recognized  that  the  French  chemist,  Moissan,  was  able  to 
prepare  the  free  element  by  the  electrolysis  of  anhydrous 
hydrofluoric  acid,  HF,  at  a  low  temperature.  He  added  a 
small  amount  of  potassium  fluoride,  KF,  to  the  acid  because 
the  pure  acid  is  almost  a  non-conductor  of  electricity.  By 
electrolyzing  the  solution  of  potassium  fluoride  in  anhydrous 
hydrofluoric  acid  in  a  U-tube  of  platinum  or  of  copper  at  a 
temperature  of  —  80°  or  below,  hydrogen  is  liberated  at  the 
cathode  and  fluorine  at  the  anode. 

At  ordinary  temperatures  fluorine  is  a  gas  of  a  light, 
greenish  yellow  color,  similar  to  the  color  of  diluted  chlorine. 
It  combines  directly  with  almost  all  other  elements  except 
oxygen.  Amorphous  carbon,  silicon,  phosphorus  and  some 
other  elements  take  fire  and  burn  when  brought  in  contact 
with  fluorine.  In  each  case  a  fluoride  is  formed,  carbon 
tetrafluoride,  CF4,  silicon  tetrafluoride,  SiF4,  and  phos- 
phorus pentafluoride,  PF5. 

With  the  exception  of  the  elements  of  the  argon  family, 
which  do  not  combine  with  other  elements  at  all,  fluorine 
is  the  only  element  which  does  not  combine  with  oxygen. 

Compounds  of  Fluorine.  Etching  Glass. — Hydrofluoric 
acid,  HF  or  H2F2,  is  prepared  by  the  action  of  concentrated 
sulfuric  acid  on  calcium  fluoride  on  the  same  principle  as 
the  preparation  of  hydrochloric  acid  from  salt: 

NaCl    +  H2SO4±=>NaHSO4  +  HC1 
CaF2  +  H2SO4<=»CaSOi  +  2HF 


In  each  case  the  acid,  hydrochloric  acid  or  hydrofluoric 
acid,  escapes  from  the  mixture  because  it  is  a  gas  and  the 
disturbance  of  the  equilibrium  causes  the  reaction  to  go 
always  in  the  direction  to  form  more  of  the  compound 
which  escapes. 

One  of  the  most  interesting  and  important  properties 
of  hydrofluoric  acid  is  its  effect  on  glass.  Ordinary  glass 


COMPARISON  97 

consists  mainly  of  a  mixture  of  silicates  of  calcium  and  so- 
dium. When  hydrofluoric  acid  gas  comes  in  contact  with 
it  the  fluorine  unites  both  with  the  silicon  and  with  the 
metals,  forming  silicon  tetrafluoride,  SiF4,  calcium  fluoride, 
*  CaF2,  and  sodium  fluoride,  NaF.  The  hydrogen  of  the 
acid  combines  with  the  oxygen  of  the  silicate  to  form  water. 
The  silicon  tetrafluoride,  SiF4,  is  a  gas  and  escapes.  Hydro- 
fluoric acid  does  not  affect  wax  or  paraffin.  If  a  piece  of 
glass  is  covered  with  paraffin  and  a  portion  of  the  surface 
of  the  glass  is  exposed  by  drawing  lines  or  figures  through 
the  wax,  on  exposing  the  object  to  the  action  of  hydrofluoric 
acid  gas  the  parts  of  the  glass  uncovered  will  be  etched. 
In  this  manner  the  scales  of  thermometers  and  of  other 
instruments  are  made.  In  some  cases  solutions  or  pasty 
mixtures  containing  hydrofluoric  acid  are  used  in  place  of  the 
gas. 

Comparison  of  the  Elements  of  the  Halogen  Family. — 
The  colors  of  the  halogens  grow  deeper  with  increasing 
atomic  weight.  Fluorine  is  light  greenish  yellow;  chlorine 
is  similar,  but  of  a  deeper  shade;  bromine  vapor  is  reddish 
brown,  and  iodine  vapor  is  violet.  Solid  iodine  is  black. 

At  room  temperature  fluorine  and  chlorine  are  gases; 
bromine  is  a  liquid,  iodine  is  a  solid. 

All  of  the  family  combine  with  hydrogen  and  the  metals. 
The  following  compounds  may  be  taken  as  illustrations: 

HF  HC1  HBr  HI 

NaF  NaCl  NaBr  Nal 

CaF2  CaClo  CaBr2  CaI2 

A1F3  A1C13  AlBr3  A1I3 

In  these  compounds  with  hydrogen  and  the  metals  the 
affinity  of  the  halogens  decreases  with  increasing  atomic 
weight.  Hydriodic  acid,  HI,  undergoes  considerable  de- 
composition if  heated  to  300°  or  400°,  hydrobromic  acid 
is  much  more  stable  at  that  temperature,  while  hydro- 


98  THE  HALOGEN  FAMILY 

chloric  and  hydrofluoric  acids  undergo  scarcely  any  de- 
composition, even  at  1000°.  Any  element  of  the  family 
will  take  hydrogen  or  a  metal  away  from  any  other  member 
of  the  family  having  a  higher  atomic  weight.  Fluorine 
will  liberate  chlorine  from  hydrochloric  acid,  bromine  from 
hydrobromic  acid,  or  iodine  from  hydriodic  acid.  Bromine 
will  liberate  iodine  from  hydriodic  acid  or  from  potassium 
iodide,  KI. 

In  the  compounds  with  oxygen  the  affinity  of  the  elements 
increases  with  increasing  atomic  weight.  Fluorine  does 
not  combine  with  oxygen;  chlorine  monoxide,  CljO,  and 
chlorine  peroxide,  C1O2,  explode  when  heated,  while  iodine 
pentoxide,  I2O5,  is  a  comparatively  stable  solid. 

The  oxygen  acids  increase  in  stability  with  an  increase  in 
the  amount  of  oxygen.  Perchloric  acid,  HC104,  is  much 
more  stable  than  chloric  acid,  HC103,  and  chloric  acid  is 
more  stable  than  hypochlorous  acid. 

SUMMARY 

There  are  several  groups  of  closely  related  elements. 

The  halogen  family  (Group  VII)  consists  of  fluorine, 
chlorine,  bromine  and  iodine. 

Bromides  are  found  in  brines,  with  chlorides. 

Bromine  forms  bromides,  hypobromites  and  bromates. 

Sodium  bromide  is  much  used  in  medicine ;  silver  bromide 
in  photography. 

Iodides  are  found  in  sea  water  and  in  the  ash  of  sea  weeds. 

Iodine  forms  iodides,  iodates  and  periodates. 

Iodine  dissolves  in  alcohol.  It  gives  a  dark  blue  color 
with  starch. 

Tincture  of  iodine,  potassium  iodide  and  an  organic 
compound  of  iodine  from  the  thyroid  gland  are  used  in 
medicine. 

Fluorine  occurs  in  calcium  fluoride,  or  fluorspar. 


EXERCISES.     HALOGENS  99 

Fluorine  is  prepared  by  the  electrolysis  of  anhydrous 
hydrofluoric  acid  containing  potassium  fluoride.  It  is 
the  most  active  of  the  non-metallic  elements. 

Hydrofluoric  acid  is  prepared  by  the  action  of  sulfuric 
acid  on  fluorspar.  It  is  used  in  etching  glass. 

The  halogens  have  a  negative  valence  (pp.  104  and  106) 
of  one  when  combined  with  hydrogen  or  the  metals.  In 
these  compounds  each  will  displace  any  of  those  of  higher 
atomic  weight. 

Fluorine  does  not  combine  with  oxygen.  The  stability  of 
the  compounds  of  the  other  halogens  with  oxygen  increases 
with  increasing  atomic  weight  and  with  increasing  amounts 
of  oxygen. 

EXERCISES 

1.  One  liter  of  hydrobromic  acid  gas  weighs  approximately 
3.62  grams.     How  many  grams  of  sodium  bromide  would  be 
required  to  prepare  22.4  liters  of  the  gas  if  it  could  be  obtained 
by  the  same  method  which  is  used  for  the  preparation  of  hydro- 
chloric acid? 

2.  How  many  grams  of  sulfuric  acid  will  be  required  for  the 
same  preparation? 

3.  Sodium    bromide,    manganese    dioxide    and    sulfuric    acid 
interact  with  each   other,   giving  bromine,   manganese   sulfate, 
MnS04,  sodium  sulfate,  Na2S04,  and  water.     How  many  grams 
of  each  compound  will  be  required  to  give  160  grams  of  bromine? 

4.  A  liter  of  hydrogen  weighs  approximately  0.09  gram.     How 
many  liters  of  hydrogen  will  be  given  by  the  action  of  an  excess  of 
hydrochloric  acid  in  the  following:  39  grams  of  potassium,  56 
grams  of  iron,  27  grams  of  aluminium,  1 18  grams  of  tin  (giving 
SnCl2),  65  grams  of  zinc? 


CHAPTER  X 
CLASSIFICATION  OF  ELEMENTS.     VALENCE 

Other  Families  of  Elements. — The  elements  of  Group  VI, 
the  sulfur  family,  with  one  exception,  have  atomic  weights  a 
little  less  than  those  of  the  chlorine  family.  The  elements  of 
the  nitrogen  family,  in  turn,  have  atomic  weights  a  little 
less  than  those  of  the  sulfur  family,  and  the  atomic  weights  of 
the  carbon  family  are  less  than  those  of  the  nitrogen  family. 
The  elements  of  the  family  of  " Noble  gases,"  which  have  no 
chemical  affinity,  have  atomic  weights  a  little  greater  than 
those  of  the  chlorine  family.  The  following  table  gives 
boron  in  addition  to  the  families,  or  groups  of  elements 
mentioned.  It  includes  all  of  the  non-metallic  elements  and 
a  few  half-metals  and  metals.  This  table  furnishes  a  very 
convenient  basis  for  remembering  the  properties  and  rela- 
tions of  these  elements  and  the  student  is  advised  to  memo- 
rize the  atomic  weights  of  the  elements  of  each  family  as 
the  successive  families  are  studied. 


Group  III 

Group  IV 

Group  V 

Group  VI 

Group  VII 

Group  O 

He  4 

Bll 

C12 

N14 

O  16 

F19 

Ne20 

A127 

Si  28 

P31 

S32 

Cl  35.5 

A  40 

Ga  69.9 

Ge  72.5 

As  75 

Se79 

Br80 

Kr83 

In  114.8 

Sn  118 

Sbl20 

Te  127.6 

1127 

Xel30 

T1204 

Pb  207 

Bi  208 

Nt222 

A  study  of  these  elements  has  made  it  very  clear  that 
there  is  a  close  .connection  between  the  properties  of  an 

100 


VALENCE  A 


element  and  its  atomic  weight.  The  table  is  part  of  a 
larger  table  in  which  all  of  the  elements  are  arranged  in 
accordance  with  their  atomic  weights.  This  larger  table, 
which  will  be  considered  later  (p.  162),  is  used  as  the  basis 
for  the  classification  of  the  elements. 

Valence. — The  atomic  weight  of  sodium  is  23,  that  of 
magnesium  is  24  and  that  of  aluminium  is  27. 1  In  accord- 
ance with  the  atomic  theory,  if  we  take  23  milligrams  of 
sodium,  24  milligrams  of  magnesium  and  27  milligrams  of 
aluminium,  we  shall  have  the  same  number  of  atoms  of  each 


FIG.    27. 


element.  If  we  allow  these  quantities  of  the  elements  to 
act  on  hydrochloric  acid,  we  shall  get  from  the  sodium  about 
11  cc.  of  hydrogen  gas,  from  the  magnesium  22  cc.  and  from 
the  aluminium  33  cc.  (Fig.  27) .  It  is  evident  from  this  that 
one  atom  of  magnesium  displaces  twice  as  much  hydrogen 
and  combines  with  twice  as  much  chlorine  as  one  atom  of 
sodium;  and  that  one  atom  of  aluminium  displaces  and 
combines  with  three  times  as  much  hydrogen  and  chlorine 

1  These  atomic  weights  are  selected  on  the  basis  of  the  specific  heats  of 
the  elements  (see  p.  266.) 


102  CLASSIFI CATION  OF  ELEMENTS 

as  the  sodium  atom.  These  relations  are  clearly  expressed 
in  the  following  equations: 

Na  +  HC1  =  NaCl  +  H 
Mg  +  2HC1  =  MgCl2  +  2H 
Al  +  3HC1  =  A1C1,  +  3H 

Similar  differences  between  the  non-metallic  elements  are 
illustrated  by  the  following  tables  of  their  compounds  with 
hydrogen : 

CH4  NH3  OH2  FH 

SiH4  PH3  SH2  C1H 

AsH3  SeH2  BrH 

SbH3  TeH2  IH 

A  study  of  these  and  other  compounds  has  led  to  the 
conclusion  that  in  chemical  compounds  each  atom  unites 
directly  with  only  a  small  number  of  other  atoms  and  that 
this  number  of  atoms  with  which  an  atom  combines  varies 
for  the  different  elements.  This  property  is  called  valence. 
Atoms,  such  as  those  of  sodium  or  chlorine,  which  combine 
with  a  single  atom  of  chlorine  or  hydrogen,  are  said  to  have  a 
valence  of  one  and  are  called  univalent.  Those  which  unite 
with  two  atoms,  such  as  the  atoms  of  magnesium  and  of 
oxygen,  have  a  valence  of  two  and  are  called  bivalent. 
Aluminium  and  phosphorus  have  a  valence  of  three  and  are 
called  trivalent.  Carbon  and  silicon  have  a  valence  of 
four  and  are  called  quadrivalent.  Oxygen  is  evidently 
bivalent  in  water,  H2O,  and  it  seems  to  be  bivalent  in  nearly 
or  quite  all  of  its  stable  compounds.  Taking  this  into  con- 
sideration in  the  series 

Na20,  MgO,  A12O3,  SiO2,  P205,  S03,  C12O7 

the  valence  of  the  elements  increases  from  one  to  seven  and 
the  elements  are  called,  successively,  univalent,  bivalent, 
trivalent,  quadrivalent,  quinquivalent,  sexivalent  and 
septivalent. 


GRAPHICAL  FORMULAS  103 

Varying  Valences. — The  valence  of  elements  is  not, 
however,  so  simple  as  the  statements  of  the  previous  para- 
graph would  seem  to  imply.  Phosphorus  has  a  valence  of 
three  in  the  compound  PH3,  but  it  forms  two  compounds, 
PC13  and  PC15,  with  chlorine.  In  the  first  of  these  it  is 
trivalent  but  in  the  second  it  is  quinquivalent. 

In  such  a  series  of  compounds  as  the  following: 

N20,  NO,  N203,  N02,  N205 
nitrogen  seems  to  vary  in  its  valence  from  one  to  five. 

While  it  is  clear  from  the  illustrations  given  that  the  same 
element  may  have  a  different  valence  in  different  compounds 
the  fundamental  idea  of  valence,  that  in  compounds  each 
atom  holds  directly  to  only  a  small,  definite  number  of 
other  atoms,  has  proved  very  useful. 

Graphical  Formulas. — The  valence  of  elements  in  com- 
pounds may  be  clearly  expressed  by  formulas  in  which  the 
symbols  of  the  elements  are  taken  to  represent  atoms,  and 
lines  are  used  to  represent  the  direct  attachments  which 
exist  between  the  atoms.  This  gives  such  formulas  as  the 
following,  which  will  be  easily  understood  without  further 
explanation. 

/Cl  /Gl 

Na  -  Cl        MgQ  Al^Cl 

XC1  XC1 

H  H 

I  / 

H  -  C  -  H        N— H        H-O-H      H-Cl 

I  \ 

H  H 


J3  ,,0 

Mg=o      \>    off       No   s^o        ;o 


104  CLASSIFICATION  OF  ELEMENTS 

Such  formulas  are  called  graphical  formulas. 

Groups  of  elements  may  also  have  a  definite  valence,  as 
is  seen  in  such  compounds  as  sodium  sulfate,  Na2S04,  and 
sodium  phosphate,  Na3PO4,  in  which  the  group  SO4  is 
bivalent  and  the  group  P04  trivalent.  This  is  more  clearly 
expressed  by  the  formulas: 


Na, 


N 
N 
Na'  Na 

In  all  such  cases  the  valences  of  the  two  parts  must  agree 
and  this  gives  valuable  assistance  in  the  writing  of  correct 
formulas. 

It  is  supposed,  of  course,  that  the  valence  of  the  group 
is  dependent,  ultimately,  on  the  valences  of  the  atoms  of 
which  it  is  composed.  This  may  be  expressed  by  such 
formulas  as: 

Na  -  Ov        ,£>  Na  -  (\ 

)S^        and     Na  -  (A?  =  O 

Na  -  O/    ^O  Na  -  O/ 

Positive  and  Negative  Valences.  —  In  the  electrolysis 
of  hydrochloric  acid,  HC1,  sodium  chloride,  NaCl,  or  of  other 
metallic  chlorides,  the  chlorine  goes  to  the  anode  or  positive 
pole  while  the  hydrogen  or  metal  goes  to  the  cathode  or 
negative  pole.  The  same  is  true  of  hydrogen  and  metallic 
compounds  of  the  other  halogens,  fluorine,  bromine  and 
iodine.  For  this  reason  hydrogen  and  the  metals  are 
called  positive  because  they  are  attracted  by  the  negative 
electrode,  and  the  halogens  are  called  negative  because 
they  are  attracted  by  the  positive  electrode.  In  the 
electrolysis  of  dilute  sulfuric  acid,  or  of  a  solution  of 
sodium  hydroxide,  hydrogen  is  liberated  at  the  cathode 
and  oxygen  at  the  anode  and  this  gives  us  a  good  reason  for 


POSITIVE  AND  NEGATIVE  VALENCES  105 

considering  oxygen  negative.  On  account  of  these  relations 
we  may  say  that  hydrogen  and  sodium  have  a  positive 
valence  of  one;  calcium  has  a  positive  valence  of  two; 
chlorine  and  bromine  have  a  negative  valence  of  one,  and 
oxygen  has  a  negative  valence  of  two. 

If  some  elements  were  always  positive  and  others  negative 
it  would  be  an  easy  matter  to  classify  elements  on  this  basis. 
But  some  elements,  such  as  nitrogen,  sulfur  and  many 
others,  combine  with  both  hydrogen  and  with  oxygen. 
We  have,  for  instance, 

<° 

N^H  and  )O 

H  < 

0 

Ammonia  Nitrogen  trioxide 

We  are  compelled  either  to  believe  that  nitrogen  has  a 
negative  valence  in  ammonia  and  a  positive  valence  in 
nitrogen  trioxide,  or  that  the  positive  or  negative  character 
of  an  element  disappears  when  elements  unite.  It  is  only 
fair  to  say  that  chemists  are  not  altogether  certain  which 
of  these  alternatives  is  true.  It  is  a  question  under  debate 
and  the  relations  are  so  complicated  that  a  satisfactory 
decision  may  not  be  reached  for  some  time.  Meanwhile 
the  distinction  between  positive  and  negative  valences  is 
one  of  increasing  importance  in  the  study  of  chemistry. 
In  the  following  table  positive  valences  are  indicated  by  the 
plus  sign  and  negative  valences  by  the  minus  sign. 

A  study  of  the  table  shows  that  positive  valences  are  much 
more  common  than  negative,  and  this  is  true  even  of  the 
non-metallic  elements.  No  element  with  an  atomic  weight 
above  130  shows  a  negative  valence.  It  is  also  doubtful 
if  any  element  of  the  first  three  groups,  except  boron, 
shows  a  negative  valence. 


106 


CLASSIFICATION  OF  ELEMENTS 


oo      -a 

°  ~t        . 
O  co"      _c 


«$ 


I  ^ 


+ 

M 

0      | 


+  +       + 


1 


iO   ^ 

i+ 


IO 

?SJ: 


? 


0 


„+ 

0 


o+ 


CO     C3    CO  ^H    + 
§  + 


g,<N 
»+ 


C(M  ^,<M  -O 

tSJ    +          02+          0 


M  "T 


o 


5? 


§0 

a 


SUMMARY.     VALENCE  107 

The  following  table  gives  the  most  common  valences  for 
each  group: 

Common  Valences 

Group  0 0 

Group  I +  1 

Group  II +  2 

Group  III -f  3 

Group  IV +4,  -  4 

Group  V +  3,  -h  5,  -  3(-  4  +  1) 

Group  VI +  2,  +  4,  +  6,  -  2(-  3  +  1) 

Group  VII +  1,  +  3,  +  5,  +  7,-l(-  2  +  1) 

Group  VIII +  2,  +  3',  +  4,  +  8 

SUMMARY 

A  very  satisfactory  classification  of  the  five  families  of 
non-metallic  elements  is  based  on  their  atomic  weights.  , 

An  atom  of  each  element,  except  those  of  the  zero  group, 
has  the  power  to  combine  directly  with  a  definite,  small 
number  of  other  atoms.  This  power  is  called  valence. 
Valences  vary  both  for  different  elements  and  for  the  same 
element  in  different  compounds. 

The  distribution  of  the  valences  in  compounds  may  be 
most  clearly  expressed  by  graphical  formulas. 

Some  valences,  as  those  of  hydrogen  and  the  metals, 
are  positive;  others,  as  those  of  the  halogens  in  the  halides, 
are  negative. 

Elements  with  positive  valences  are  much  more  common 
than  those  with  negative  valences. 

A  table  is  given  for  the  valences  of  the  elements. 

EXERCISES 

Write  graphical  formulas  for  the  following  compounds :  manga- 
nese dioxide,  manganese  chloride,  manganese  sulfate,  potassium 
hypochlorite,  potassium  chlorite,  potassium  chlorate,  potassium 
perchlorate,  potassium  permanganate.  In  the  salts  represent 


108  CLASSIFICATION  OF  ELEMENTS 

the  potassium  as  united  to  the  other  elements  through  one  oxygen 
atom. 

While  the  graphical  formulas  of  inorganic  compounds  which  are 
referred  to  here  have  considerable  value  and  very  probably  re- 
present the  real  structure  of  these  compounds,  too  much  value 
should  not  be  attached  to  them.  Especially,  it  is  important  to 
understand  that  valence  alone  is  not  a  sufficient  guide  for  the  writing 
of  graphical  or  structural  formulas.  The  two  formulas 

^° 
H-  0-0-0-0- Cl  and  H  -  0  -  Cl  =0  are  both  in  ac- 

^0 

cord  with  the  valence  of  chlorine  in  other  compounds  but  only 
one  of  these  can  be  true  and  we  are  not  altogether  certain  that 
either  is  true,  though  there  is  considerable  probability  of  the  truth 
of  the  second. 


CHAPTER  XI 
GROUP  VI:  SULFUR,  SELENIUM  AND  TELLURIUM 

Occurrence. — Chlorine  and  the  other  elements  of  the 
chlorine  family  are  not  found  free  in  nature,  chiefly  because 
of  their  strong  affinity  for  metals  which  are  common  and 
abundant.  Oxygen,  on  the  other  hand,  is  found  free  partly 
because  of  the  great  abundance  of  the  element  and  partly 
because  the  growth  of  plants  constantly  liberates  oxygen 
from  the  carbon  dioxide  of  the  air.  Sulfur,  the  second 
element  of  the  oxygen  group,  is  also  found  in  the  free  state 
largely  because  oxygen  liberates  sulfur  from  hydrogen 
sulfide  very  much  as  fluorine  would  liberate  chlorine  from 
hydrochloric  acid  or  as  chlorine  liberates  bromine  from 
bromides. 

Free  sulfur  is  found  especially  in  Sicily  and  in  Louisiana. 
For  a  long  time  the  sulfur  mines  of  Sicily  held  a  practical 
monopoly  of  the  sulfur  markets  of  the  world.  The  sulfur 
of  Louisiana  is  beneath  a  bed  of  quicksand  and  for  many 
years  after  the  deposits  were  discovered  no  one  was  able 
to  devise  a  method  for  mining  the  mineral.  Finally 
Hermann  Frasch,  of  New  York,  conceived  the  idea  that  it 
would  be  possible  to  melt  the  sulfur  by  means  of  hot  water 
under  pressure  and  he  devised  a  system  of  concentric  tubes 
by  means  of  which  hot  water  is  brought  down  on  top  of  the 
sulfur  and  the  melted  sulfur  is  brought  to  the  surface  by 
means  of  a  current  of  air.  This  process  now  furnishes  all 
of  the  free  sulfur  required  for  manufacturing  purposes  in 
the  United  States. 

In  Sicily  the  sulfur  is  mixed  with  other  minerals  from 

109 


110  SULFUR,  SELENIUM  AND  TELLURIUM 

which  it  is  separated  by  melting  it.  The  heat  is  furnished 
sometimes  by  burning  a  part  of  the  sulfur,  as  other  fuels 
are  scarce  and  expensive.  The  sulfur  is  further  purified  by 
distilling  it.  If  the  vapors  are  condensed  in  warm  chambers 
so  that  the  sulfur  remains  liquid,  it  is  cast  in  rolls  and  called 
roll  brimstone.  If  the  condensing  chambers  are  cold  so 
that  the  vapor  condenses  at  once  to  a  solid  the  product  is 
called  flowers  of  sulfur.  The  difference  between  the  for- 
mation of  roll  brimstone  and  flowers  of  sulfur  is  similar  to 
that  between  the  condensation  of  water  vapor  from  the 
clouds  as  rain  and  as  snow. 

Sulfur  is  also  found  in  nature  in  the  form  of  sulfides  and 
of  sulfates.  Several  important  ores  of  common  metals  are 
sulfides.  Among  these  may  be  mentioned  galena,  or  lead 
sulfide,  PbS,  sphalerite,  or  zinc  sulfide,  ZnS,  and  chalcopy- 
rite,  or  copper  pyrites,  CuFeS2.  Iron  pyrites  or  pyrite, 
FeS2,  is  extensively  used  as  a  source  of  sulfur  for  the  manu- 
facture of  sulfuric  acid,  and  is  to  be  considered  as  an  ore 
of  sulfur  rather  than  as  an  ore  of  iron.  The  most  important 
sulfates  are  calcium  sulfate,  found  as  the  hydrate,  CaSO4. 
2H2O,  called  gypsum  or  alabaster  and  used  in  the  manu- 
facture of  plaster  of  Paris,  and  barium  sulfate,  BaSO4,  used 
in  making  compounds  of  barium  and  as  a  substitute  for 
white  lead. 

Sulfur  is  used  extensively  in  "sulfuring  fruit"  which  is  to 
be  dried,  the  fruit  being  exposed  to  the  action  of  sulfur 
dioxide  obtained  by  burning  the  sulfur.  The  sulfur  diox- 
ide kills  microorganisms  which  would  injure  the  fruit  before 
it  is  dry  and  also  prevents  the  darkening  of  the  fruit. 

Sulfur  is  also  dissolved  in  milk  of  lime  to  give  a  "lime- 
sulfur  wash"  which  is  applied  to  vines  and  fruit  trees  to 
prevent  the  growth  of  fungi  and  other  harmful  organisms. 

Sulfur  is  a  constituent  of  the  old-fashioned  black  gun- 
powder but  is  not  used  in  the  smokeless  powders.  It  is  used 
in  the  manufacture  of  india-rubber  and  of  carbon  disul- 


FORMS  OF  SULFUR  111 

fide  and  in  some  dyes,  especially  in  sulfur  black.  It  is 
sometimes  used  in  making  sulfuric  acid,  but  iron  pyrites, 
FeS2,  is  usually  a  cheaper  source  of  sulfur  for  that  purpose. 

Forms  of  Sulfur. — There  are  two  allotropic  forms  of 
oxygen,  ordinary  oxygen  and  ozone.  Sulfur  gives  three 
allotropic,  solid  forms. 

The  sulfur  found  in  nature  crystallizes  in  rhombic  crystals 
which  melt  at  114.5°.  Similar  crystals  can  be  prepared 
by  dissolving  sulfur  in  carbon  disulfide  and  allowing  the 
solution  to  evaporate. 

If  sulfur  is  heated  till  it  melts  and  then  is  allowed  to  cool 
slowly  it  crystallizes  in  long  needles  which  can  be  obtained 
by  breaking  the  crust  on  the  surface,  when  the  mass  is 
partially  solidified,  and  pouring  out  the  liquid  portion. 
These  needles  are  nearly  transparent  at  first  and  melt 
at  119°.  If  they  are  allowed  to  cool  and  stand  for  some 
time  they  become  opaque  and  the  melting  point  changes 
to  114.5°,  showing  that  they  have  changed  to  microscopic 
crystals  of  the  rhombic  variety.  It  is  evident  from  this 
that  the  rhombic  crystals  are  the  stable  form  at  ordinary 
temperatures  while  the  needles  are  more  stable  at  a  tempera- 
ture near  the  melting  point. 

When  sulfur  is  heated  above  its  melting  point,  at  160° 
it  suddenly  becomes  thick  and  viscous.  If  it  is  cooled 
slowly  it  changes  back  as  it  cools  to  the  ordinary  forms, 
but  if  it  is  cooled  quickly  it  may  be  obtained  in  a  form  which 
is  at  first  plastic  like  india-rubber  but  which  may  solidify 
without  crystallizing.  This  is  called  amorphous  sulfur, 
meaning  that  it  has  no  crystalline  form  and  it  is  to  be  con- 
sidered as  a  supercooled  liquid  rather  than  as  an  ordinary 
solid.  This  amorphous  sulfur  is  insoluble  in  carbon 
disulfide. 

Liquid  and  Gaseous  Sulfur. — At  a  temperature  slightly 
above  its  melting  point  sulfur  is  a  pale,  yellow,  mobile 
liquid.  At  160°,  as  stated  above,  it  suddenly  becomes 


112  SULFUR,  SELENIUM  AND  TELLURIUM 

very  viscous  and  the  character  of  the  substance  obtained 
by  cooling  this  liquid  shows  that  the  sulfur  changes  to  an 
allotropic  form  at  that  temperature.  When  heated  still 
hotter  the  liquid  becomes  more  mobile  again,  but  remains 
of  the  same  dark  color  which  is  characteristic  of  the  viscous 
liquid.  At  445°  the  sulfur  boils.  If  care  is  taken  to  keep 
the  walls  of  the  tube  or  flask  above  the  boiling  liquid  at 
a  temperature  above  the  boiling  point  it  can  be  seen  that 
the  vapor  is  of  a  light  amber  color. 

Hydrogen  Sulfide. — Waters  from  artesian  wells  often 
contain  hydrogen  sulfide  in  solution  and  such  waters  are 
known  'as  "sulfur"  waters.  The  hydrogen  sulfide  gives 
to  the  waters  a  disagreeable  odor  and  taste.  Hydrogen 
sulfide  is  one  of  the  substances  which  gives  the  offensive 
odor  of  decayed  eggs  and  it  is  usually  present  in  sewer  gas. 
It  is  prepared  by  the  action  of  hydrochloric  or  sulfuric 
acid  on  ferrous  sulfide,  FeS,  just  as  water  is  formed  when 
an  acid  acts  on  zinc  oxide,  ZnO: 

ZnO  +  2HC1  =  ZnCl2  +  H2O 
FeS  +  2HC1  =  FeCl2  +  H2S 

Ferrous  Ferrous 

.  sulfide  chloride 

Hydrogen  sulfide  burns  in  air  to  sulfur  dioxide,  SO2, 
and  water.  If  a  test-tube  or  cylinder  filled  with  the  gas 
is  ignited  at  the  mouth  so  that  the  gas  burns  with  an  in- 
sufficient supply  of  air,  the  hydrogen  burns  to  water  while 
a  part  of  the  sulfur  is  separated  in  the  free  state. 

Sulfides.  Groups  of  Metals  for  Analysis. — Many  of  the 
metals  combine  directly  with  sulfur  when  heated  with  it. 
Many  of  the  sulfides  of  metals  are  insoluble  in  water 
and  the  addition  of  hydrogen  sulfide  to  solutions  of  salts 
of  these  metals  will  cause  the  sulfides  to  separate  as  pre- 
cipitates. Thus  if  hydrogen  sulfide  is  passed  into  an  acid 
solution  containing  lead,  arsenic  or  antimony,  a  precipitate 


GROUPS  OF  QUALITATIVE  ANALYSIS  113 

of  lead  sulfide,  PbS,  arsenic  sulfide,  As2S3,  or  antimony 
sulfide,  Sb2S3,  will  separate.  From  a  neutral  or  alkaline 
solution  containing  zinc  or  iron,  zinc  sulfide,  ZnS,  or  ferrous 
sulfide,  FeS,  will  separate.  This  conduct  of  solutions  of  the 
metallic  salts  furnishes  a  ready  means  of  separating  the 
metals  into  three  groups  for  purposes  of  analysis.  The 
first  group  contains  lead,  arsenic,  antimony  and  other 
metals  whose  sulfides  are  so  insoluble  that  they  may  be 
precipitated  from  acid  solutions  by  .hydrogen  sulfide. 
The  second  group  consists  of  metals  whose  sulfides  are 
slightly  more  soluble  and  for  that  reason  are  not  precipitated 
from  acid  solutions  but  they  are  precipitated  from  alkaline 
solutions. , 

If  an  acid  solution  containing  metals  of  all  three  groups 
is  treated  with  hydrogen  sulfide,  the  sulfides  of  the  first 
group  are  precipitated  and  by  filtering  off  and  washing 
the  precipitates  the  metals  of  that  group  may  be  separated 
from  the  metals  of  the  second  and  third  groups.  By  adding 
ammonia  to  the  filtrate,  that  is,  to  the  solution  containing 
metals  of  the  second  and  third  groups  which  runs  through 
the  filter,  the  second  group  will  be  precipitated  .and  this 
precipitate  may  be  filtered  off  and  separated  from  the 
metals  of  the  third  group.1 

Formation  of  Precipitates.  —  In  a  previous  chapter 
(p.  65)  the  formation  of  a  precipitate  of  silver  chloride, 
AgCl,  has  been  explained  as  due  to  the  presence  of  ions 
which  readily  separate  from  each  other  or  combine  to  form 
new  compounds  in  solution.  An  extension  of  the  same 
principle  helps  us  to  understand  the  precipitation  of  sul- 
fides. For  a  better  understanding  of  what  occurs  two  other 
principles  are  needed.  These  are: 

First,  that  the  formation  of  ions  and  the  formation  of 

1  In  the  classification  adopted  in  most  books  on  qualitative  analysis 
the  metals  of  Group  I  as  given  here  are  separated  into  two  groups  called 
I  and  II;  the  metals  of  Group  II  above  are  given  as  Group  III  and  Group 
III  above  is  divided  iato  Groups  IV  and  V. 
8 


114  SULFUR,  SELENIUM  AND  TELLURIUM 

precipitates  and  other  compounds  in  solutions  are  reversible 
reactions. 

Second,  that,  as  with  all  other  reversible  reactions,  the 
increase  of  one  of  the  substances  taking  part  tends  to  in- 
crease the  amount  of  the  substances  formed  from  it  and  a 
decrease  of  the  substance  tends  to  decrease  the  amount 
formed. 

The  precipitation  of  lead  sulfide  from  a  solution  of  lead 
nitrate,  Pb(NOs)2,  depends  on  the  reversible  reactions  rep- 
resented by  the  following  equations:1 

Pb(NO3)2  <=*  Pb++  +  N03~  +  NOr 
H2S  <=±  S-  +  H+  +  H+ 
H+  +  N03-  *±  HN03 
Pb++  +  S-  *±  PbS 

Combining  these  we  have: 

Pb(N03)2  +  H2S  =  PbS  +  2HN03 
4 

Nitric  acid  ionizes  to  a  much  greater  extent  than  hydrogen 
sulfide,  or,  in  other  words,  a  solution  of  nitric  acid  contains 
many  hydrogen  ions  while  one  of  hydrogen  sulfide  contains 
relatively  few  of  such  ions.  If  we  add  nitric  acid  or  any 
other  acid  which  gives  a  large  number  of  hydrogen  ions  the 
reversible  reaction  represented  by  the  equation : 

H2S  <=±  S-  +  H+  +  H+ 

must  be  driven  to  the  left  by  the  increase  in  one  of  the  con- 
stituents, H+,  of  the  right-hand  side  of  the  equation.  This 
will  cause  a  decrease  in  the  number  of  sulfur  ions,  S=,  and 
this  may  go  so  far  as  to  cause  the  reaction  represented  by 
the  equation: 

Pb++  +  S-  <=±  PbS 

1  For  the  sake  of  simplicity  the  formation  of  intermediate  ions  such  as 
HS~  and  Pb+  SH  and  the  hydrolysis  of  Na2S  to  NaHS  and  NaOH  are 
not  considered  here.  See  Noyes,  Textbook  of  Chemistry,  pp.  169,  170 


PRECIPITATION.     SULFUR  DIOXIDE  115 

also  to  go  to  the  left  and  the  lead  sulfide  would  dissolve. 
It  follows  from  this  that  while  lead  sulfide  is  precipitated 
from  a  solution  containing  a  dilute  acid  it  will  dissolve  in  a 
more  concentrated  acid  and  the  precipitate  will  not  be 
formed  in  the  presence  of  much  free  acid. 

The  precipitation  of  ferrous  sulfide  from  an  alkaline  solu- 
tion depends  on  the  fact  that  in  the  reaction  represented  by 
the  equation: 

Na2S  ?=*  S-  +  Na+  +  Na+ 

the  equilibrium  is  much  farther  to  the  right  and  there  are 
many  more  sulfur  ions,  S=,  present  than  in  the  correspond- 
ing ionization  of  hydrogen  sulfide.  The  larger  number  of 
sulfur  ions  causes  the  formation  of  ferrous  sulfide : 

FeCl2  <=*  Fe++  +  Cl~  +  Cl~ 
Fe++  +  S-  <F±  FeS 

In  all  reactions  when  a  precipitate  is  formed  the  compound 
precipitated  is  removed  from  the  sphere  of  action  and  this 
drives  the  reaction  in  the  direction  toward  its  formation. 

Hydrosulfides. — When  a  hydroxide  of  a  metal  which  is 
not  precipitated  by  hydrogen  sulfide  is  treated  with  the  gas 
a  hydrosulfide  is  formed : 

NaOH  +  HSH  =  NaSH  +  HOH 

The  name  hydrosulfide  refers,  of  course,  to  a  compound 
containing  hydrogen  and  sulfur,  just  as  a  hydroxide  con- 
tains hydrogen  and  oxygen. 

Sulfur  Dioxide. — Sulfur  burns  in  air  or  in  oxygen  with  a 
blue  flame,  giving  a  gas  with  a  suffocating  odor.  This  has 
the  composition  represented  by  the  formula  S02  and  is 
called  sulfur  dioxide.  The  same  odor  is  sometimes  ob- 
served from  the  cinders  of  coal  containing  pyrites.  Sulfur 
dioxide  is  a  good  germicide  and  its  use  in  "sulfuring"  fruit 
has  been  spoken  of.  It  was  formerly  used  to  disinfect 


116  SULFUR,  SELENIUM  AND  TELLURIUM 

rooms,  clothing  and  articles  which  had  been  infected  with 
bacteria  which  cause  infectious  diseases,  but  its  use  for  these 
purposes  has  been  almost  entirely  replaced  by  the  use  of 
formaldehyde,  which  is  even  more  effective  and  which  does 
not  bleach  fabrics  and  discolor  metallic  objects  as  sulfur 
dioxide  does. 

Sulfur  dioxide  is  used  for  bleaching,  especially  for  bleach- 
ing straw  hats.  It  produces  its  effect  by  combining  with  the 
coloring  matter  of  the  straw  to  produce  a  colorless  com- 
pound. These  compounds  are  not  very  stable  and  they 
are  decomposed  by  the  action  of  air  and  light,  restoring 
the  original  color  of  the  straw.  The  action  in  this  regard  is 
quite  different  from  that  of  chlorine,  which  permanently 
destroys  the  coloring  matters  which  it  attacks.  Chlorine  is 
more  liable  to  weaken  the  fiber  of  the  material  which  is 
treated,  however. 

Sulfur  dioxide  may  be  prepared  by  warming  copper  with 
concentrated  sulfuric  acid.  We  may  suppose  that  the 
acid  oxidizes  the  copper  to  cupric  oxide,  CuO,  and  is  itself 
reduced  to  sulfurous  acid,  H2SO3,  and  that  the  latter  de- 
composes into  sulfur  dioxide  and  water.  The  copper  oxide 
also  reacts  with  more  sulfuric  aeidr  giving  copper  sulfater 
CuS04: 

Cu  +  H2SO  =  [CuO];  -f  H2S03 

H2SO3  =  SO2  -f  H2O 
.[CuO]  +  H2SO4  =  CuSO4  -f  H2O 

Combining: 

Cu  +  2H2SO4  =  CuSO4  +  $O2  +  2H20 

The  copper  oxide,  CuO,  is  enclosed  in  brackets  because 
we  have  no  positive  evidence  that  it  is  formed  as  an  inter- 
mediate compound.  We  only  know  that  the  reaction  takes 
place  in  accordance  with  the  last  equation  and  that  the 
supposition  that  copper  oxide  is  io-raied!  temporarily  is  in 


SULFITES.     DIBASIC  ACIDS  117 

accordance  with  the  facts  that  sulfuric  acid  is  reduced  and 
that  the  copper  is  oxidized  in  the  reaction. 

Sulfurous  Acid.  Sulfites. — Water  dissolves  a  certain 
amount  of  sulfur  dioxide.  Part  of  the  gas  combines  with 
the  water  to  form  sulfurous  acid,  H2SO3,  as  is  shown  by  the 
acid  reaction  of  the  solution  with  litmus  paper  and  by  the 
neutralization  of  bases  by  the  solution.  By  passing  sulfur 
dioxide  into  a  solution  of  sodium  hydroxide,  NaOH,  or 
milk  of  lime,  Ca(OH)2,  sodium  sulfite,  Na2SO3,  or  calcium 
sulfite,  CaS03,  may  be  prepared.  The  same  salts  may  be 
prepared  by  the  action  of  sulfur  dioxide  and  water  on  so- 
dium carbonate,  Na2CO3,  or  calcium  carbonate,  CaCO3 
because  sulfurous  acid  dissociates  (p.  150)  less  easily  into 
sulfur  dioxide  and  water  than  carbonic  acid  dissociates  to 
carbon  dioxide  and  water. 

Acid  Sulfites.  Dibasic  Acids. — Sulfurous  acid,  H2SO3, 
has  two  atoms  of  hydrogen  in  its  molecules  which  may  be 
replaced  by  metals  and  it  not  only  forms  salts,  such  as  sodium 
sulfite,  Na2SO3,  and  calcium  sulfite,  CaS03,  in  which  all  of 
the  hydrogen  Tias  been  replaced,  but  it  also  forms  salts  in 
which  only  one-half  of  the  hydrogen  has  been  replaced. 
Acid  sodium  sulfite,  NaHS03,  and  acid  calcium  sulfite, 
CaH2(SO3)2,  are  salts  of  this  type.  Such  salts  may  be 
prepared  by  adding  to  a  solution  of  sulfur  dioxide  enough  of 
the  base  to  neutralize  only  one-half  of  the  acid: 

H2S03  +  NaOH  =  NaHSO3  +  H2O 

The  salts  may  also  be  prepared  by  passing  sulfur  dioxide 
into  a  solution  of  the  base  till  the  odor  of  sulfur  dioxide 
becomes  apparent. 

2NaOH  +  SO2  =  Na2SO3  +  H2O 
Na2SO3  +  SO2  +  H2O  =  2NaHSO3 

Note. — Similar  reactions  should  be  written  for  the  prepa- 
ration of  the  calcium  salt. 


118  SULFUR,  SELENIUM  AND  TELLURIUM 

Acid  sodium  sulfite  is  a  germicide  and  it  has  sometimes 
been  added  to  cider  or  wine  to  stop  the  fermentation.  The 
acid  calcium  salt  is  used  in  paper  making  to  dissolve  un- 
desirable substances  from  the  fiber  of  the  wood  which  is 
used  to  make  paper  pulp. 

The  most  convenient  laboratory  method  for  the  prepara- 
tion of  sulfur  dioxide  is  to  drop  concentrated  sulfuric  acid 
into  a  40  per  cent  solution  of  acid  sodium  sulfite.  Write 
the  equation  for  the  reaction. 

Acids  like  sulfurous  acid,  which  form  two  classes  of  salts 
(here  NaHSO3  and  Na2SO3),  in  one  of  which  only  half  of  the 
hydrogen  is  replaced,  are  called  dibasic.  In  contrast  with 
these,  hydrochloric  acid,  HC1,  and  nitric  acid,  HNO3, 
are  monobasic.  Phosphoric  acid,  H3PO4,  which  forms  the 
salts  Na3P04,  Na2HP04  and  NaH2PO4,  is  tribasic.  A  salt 
in  which  a  metal  has  replaced  all  of  the  replaceable  hydro- 
gen of  an  acid  is  called  a  normal  salt.  Normal  salts  are  not 
always  neutral  (see  p.  77). 

Sulfur  Trioxide  or  Sulfuric  Anhydride.  Catalysis. — 
Sulfur  dioxide  will  combine  directly  with  more  oxygen  to 
form  sulfur  trioxide,  S03,  but  the  reaction  is  a  reversible 
one: 

2SO2  +  O2  ±*  2SO3 

At  ordinary  temperatures  the  equilibrium  in  this  reac- 
tion is  far  toward  the  right,  that  is,  when  the  equilibrium  is 
reached  a  large  part  of  the  sulfur  dioxide  and  oxygen  will 
have  combined  to  form  sulfur  trioxide  and  only  a  small 
part  will  remain  uncombined.  But  the  combination  takes 
place  so  slowly  that  the  equilibrium  is  reached  only  after  a 
very  long  time  and  the  reaction  could  not  be  made  commer- 
cially profitable. 

At  a  high  temperature  this  combination  takes  place 
rapidly  but  the  equilibrium  shifts  more  and  more  to  the  left, 
and  when  equilibrium  is  reached  a  large  part  of  the  sulfur 


SULFUR  TRIOXIDE  HO 

dioxide  and  oxygen  remain  uncombined.  It  is  impossible 
to  find  any  temperature  at  which  the  combination  is  both 
rapid  enough  and  complete  enough  to  be  used  as  a  commer- 
cial process. 

The  practical  solution  of  a  difficulty  of  this  kind  is  to  find, 
if  possible,  some  catalytic  agent  (p.  9)  which  will  increase 
the  speed  of  the  reaction  so  that  it  may  be  carried  on  rapidly 
at  a  temperature  low  enough  so  that  the  equilibrium  is  still 
favorable  to  the  production  of  the  compound  desired .  It  has 
been  known  for  more  than  seventy  years  that  finely  divided 
platinum  is  a  suitable  catalytic  agent  to  promote  the  union 
of  sulfur  dioxide  and  oxygen,  but  many  difficulties  were  met 
in  the  practical  details  of  the  process  and  it  was  not  till 
about  1900  that  these  were  so  far  overcome  as  to  make  the 
"contact  process"  for  the  manufacture  of  sulfur  trioxide 
and  sulfuric  acid  successful  on  a  large  scale.  At  the  present 
time  large  quantities  of  these  compounds  are  manufactured 
by  this  process. 

Properties  of  Sulfur  Trioxide. — Pure  sulfur  trioxide  is  a 
volatile  liquid  which  boils  at  46°.  A  trace  of  moisture 
causes  it  to  unite  with  itself  to  form  a  solid  compound, 
S2O6,  which  is  called  a  polymeric  form.  A  polymer  is  a 
compound  formed  by  the  union  of  two  or  more  molecules 
of  the  same  kind.  It  has  exactly  the  same  proportions  of 
the  elements  as  the  original  compound.  We  may  suppose 
that  sulfur  trioxide  and  its  polymer  have  some  such  struc- 
tures as  the  following: 

,0  0%  /Qv       .0 

0-<  and  X    )< 

^O  O^    XCK       X) 

The  preparation  of  sulfur  trioxide  on  a  small  scale  may 
be  carried  out  with  the  apparatus  shown  in  Fig.  28.  A 
mixture  of  sulfur  dioxide  and  oxygen,  dried  by  passing  the 
gases  through  bottles  containing  concentrated  sulfuric 


120 


SULFUR,  SELENIUM  AND  TELLURIUM 


acid,  is  passed  through  a  tube  containing  platinized  asbestos. 
The  tube  is  warmed  gently  and  at  a  temperature  of  300° 
to  400°  the  combination  takes  place  rapidly. 

Sulfur  trioxide  hisses  like  hot  iron  when  thrown  into  water. 
It  combines  with  the  water  to  form  sulfuric  acid.  Because 
of  this  relation  to  sulfuric  acid  it  is  often  called  sulfuric 
anhydride,  meaning  sulfuric  acid  without  water  or  from 
which  the  water  has  been  removed. 


FIG.    28. 

Sulfur  trioxide  fumes  strongly  in  the  air  because  it  com- 
bines with  the  moisture  of  the  air  to  form  sulfuric  acid. 

Effects  of  Temperature  upon  Reversible  Reactions. 
Principle  of  van't  Hoff-Le  Chatelier. — It  has  been  pointed 
out  that  an  increase  of  temperature  hastens  both  the  com- 
bination of  sulfur  dioxide  and  oxygen  and  the  decomposi- 
tion of  sulfur  trioxide,  but  that  the  rate  of  decomposition 
increases  more  rapidly  than  the  rate  of  combination.  In 
general  it  has  been  found  that  the  speed  of  any  reaction  is 
doubled  by  an  increase  of  10°  in  the  temperature.  For  an 
increase  of  20°  the  speed  is  four  times  as  great,  for  30°  eight 
times  as  great,  and  so  on.  But  the  fact  that  the  two  direc- 
tions of  a  reversible  reaction  are  not  hastened  at  the  same 
rate  shows  that  the  statement  just  given  is  only  a  rough, 
experimental  rule. 

Heat  is  generated  when  sulfur  dioxide  and  oxygen  unite. 
It  follows  that  heat  is  absorbed  when  sulfur  trioxide  is 
decomposed.  The  effect  of  temperature  on  the  equilibrium 


LEAD  CHAMBER  PROCESS  121 

is  intimately  connected  with  these  facts.  An  increase  in 
temperature  always  opposes  that  direction  of  a  reversible 
reaction  which  causes  an  evolution  of  heat.  To  take  another 
illustration:  oxygen  and  hydrogen  combine  with  an  evolu- 
tion of  heat.  At  high  temperatures  an  increase  in  tempera- 
ture causes  more  of  the  oxygen  and  hydrogen  to  remain 
uncombined  or,  what  is  the  same  thing,  causes  more  of  the 
water  to  decompose. 

An  increase  in  temperature  opposes  the  combination  of 
the  elements  because  their  combination  generates  heat  and 
the  heat  generated  tends  to  decompose  the  water  formed. 

The  effect  of  an  increase  in  temperature  upon  the  com- 
bination of  sulfur  dioxide  and  oxygen  is  an  illustration  of 
the  principle  of  van't  Hoff-Le  Chatelier,  which  is  that 
every  force  applied  to  a  system  in  equilibrium  produces  a 
change  which  tends  to  resist  the  force  applied.  If  heat  is 
generated  by  the  combination  of  two  substances  the 
application  of  heat  tends  to  decrease  the  amount  of  com- 
bination because  heat  causes  the  decomposition  of  a  part 
of  the  product  formed. 

If  we  could  understand  fully  the  mechanism  of  all  physical 
changes  it  seems  likely  that  we  should  find  that  the  principle 
has  its  foundation  in  Newton's  law  that  for  every  action 
there  is  an  equal  and  opposite  reaction. 

Lead  Chamber  Process  for  Manufacturing  Sulfuric 
Acid. — Many  different  oxidizing  agents  will  oxidize  sul- 
furous  acid,  or  sulfur  dioxide  and  water,  to  sulfuric  acid. 
One  of  these,  nitrogen  dioxide,  NO2,  is  reduced  in  the 
process  to  nitric  oxide,  NO,  a  compound  which  combines 
directly  and  rapidly  with  the  oxygen  of  the  air  to  form  the 
original  compound,  nitrogen  dioxide.  The  reactions  can 
be  expressed  most  simply  by  the  following  equations: 

N02  +  S02  +  H20  =  H2S04  +  NO 
2NO   +     02  =2NO2 


122  SULFUR,  SELENIUM  AND  TELLURIUM 

In  the  light  of  these  equations  the  nitrogen  dioxide  is 
merely  a  catalytic  agent  which  enables  sulfur  dioxide  and 
water  to  combine  rapidly  with  the  oxygen  of  the  air  to 
form  sulfuric  acid.  This  method  of  making  sulfuric  acid 
has  been  in  use  for  one  hundred  and  seventy  years  and  is 
known  as  the  "  chamber  process." 

The  manufacture  is  carried  out  in  large  chambers  built 
of  sheet  lead.  The  sheets  of  lead  are  melted  together  at 
the  edges  by  means  of  the  oxyhydrogen  blowpipe,  a  process 
known  as  "lead  burning."  and  the  chamber  of  lead  is 
supported  by  a.  wooden  framework  on  the  outside. 

The  sulfur  dioxide  for  the  process  is  usually  made  by 
burning  iron  pyrites,  FeS2,  in  a  furnace  especially  con- 
structed for  the  purpose: 

2FeS2  +  HO    =  4SO2  +  Fe2O3 
or 

4FeS2  +  11O2    =  8SO2  +  2Fe2O3 

To  furnish  the  nitrogen  dioxide  nitric  acid,  HN03,  is 
introduced.  This  reacts  at  first  with  the  sulfur  dioxide 
and  water,  giving  sulfuric  acid  and  nitric  oxide: 

2HNO3  +  3SO2  +  2H20  =  3H2SO4  +  2NO1 

Oxygen  is  furnished  by  the  introduction  of  air,  and  water 
is  introduced  in  the  form  of  a  spray. 

Since  air  is  forced  into  the  chambers  the  nitrogen  of  the 
air  must  constantly  escape  at  the  end  and  this  carries  with 
it  the  larger  part  of  the  nitrogen  dioxide.  To  prevent  the 
loss  of  this  the  gases  which  escape  from  the  chamber  are 
passed  through  a  tower  where  they  are  brought  into  contact 
with  concentrated  sulfuric  acid  running  slowly  over  pieces 
of  coke.  The  sulfuric  acid  absorbs  the  nitrogen  dioxide 

1  This  should  not  be  learned  by  rote  but  should  be  written  on  the  follow- 
ing principles:  (1)  2HNO3  will  give  2NO  +  H2O  +  3O.  (2)  Three  atoms 
of  oxygen  will  oxidize  three  molecules  of  sulfur  dioxide  to  the  trioxide. 
The  rest  follows  from  this  beginning. 


CONCENTRATION  OF  SULFURIC  ACID 


123 


but  the  gas  can  be  recovered  by  diluting  the  acid  and  bring- 
ing it  into  contact  with  the  sulfur  dioxide  as  it  comes  from 
the  pyrite  burners.  The  tower  in  which  the  nitrogen  dioxide 
is  absorbed  is  called  the  Gay-Lussac  tower  and  that  in 
which  it  is  taken  up  again  from  the  diluted  acid  is  called  the 
Glover  tower.  The  general  arrangement  of  the  whole 
apparatus  is  shown  in  Fig.  29. 


Leaden  Chambers 
FIG.    29. 


Gay-Lussac 
Tower 


Theoretically,  a  small  quantity  if  nitric  acid  should  con- 
vert an  unlimited  amount  of  sulfur  dioxide  into  sulfuric 
acid.  Practically,  there  are  unavoidable  losses  of  the 
oxides  of  nitrogen  and  in  a  well-conducted  establishment 
from  25  to  40  pounds  of  sodium  nitrate,  NaN03,  are  re- 
quired to  furnish  the  nitric  acid  necessary  for  the  manufac- 
ture of  a  ton  of  sulfuric  acid. 

Concentration  of  Sulfuric  Acid.  Properties. — The  acid 
produced  by  the  chamber  process  contains  30  to  40  per 
cent  of  water.  For  shipment  and  for  many  purposes  for 
which  the  acid  is  used  a  product  which  is  nearly  free  from 
water  is  preferred.  When  the  chamber  acid  is  heated, 
water  escapes  at  first  with  very  little  acid  and  it  is  possible 
to  carry  the  acid  up  to  a  strength  of  96  to  98  per  cent  by  sim- 
ple evaporation.  The  evaporation  is  carried  out  at  first  in 


124  SULFUR,  SELENIUM  AND  TELLURIUM 

lead  pans  which  are  scarcely  attacked  by  the  dilute  acid, 
but  as  the  acid  becomes  more  concentrated  the  lead  begins 
to  dissolve  and  the  further  concentration  up  to  93  to  95 
per  cent  is  carried  out  in  glass,  platinum  or  iron.  This 
gives  the  concentrated  commercial  acid,  which  is  usually 
slightly  brown  from  the  effect  of  dust  and  organic  matter 
and  which  contains  arsenic,  lead,  iron  and  other  impurities 
that  do  not  interfere  with  its  use  for  most  purposes.  A  pure 
and  stronger  acid  is  obtained  by  distillation.  The  concen- 
trated acid  boils  at  338°  and  contains  only  about  two  per 
cent  of  water. 

Considerable  heat  is  generated  when  water  is  mixed  with 
concentrated  sulfuric  acid.  The  concentrated  acid  has 
a  specific  gravity  of  L84  and  it  should  always  be  diluted 
by  pouring  ihe  acid  into  water  and  not  vice  versa.  If  water 
is  poured  upon  the  acid  so  much  heat  may  be  produced  at 
the  junction  of  the  two  liquids  as  to  create  an  explosion  by 
the  steam  which  is  generated. 

The  affinity  of  sulfuric  acid  for  water  is  so  great  that  it 
abstracts  the  elements  of  water  from  such  substances  as 
sugar  or  wood,  charring  them. 

Concentrated  sulfuric  acid  produces  very  painful  wounds. 
In  cases  of  accident  the  acid  should  be  removed  as  quickly 
as  possible  by  the  copious  application  of  water  and  after- 
ward moist  sodium  bicarbonate  may  be  applied  to  the 
wound. 

Because  of  its  high  boiling  point  and  low  price,  sulfuric 
acid  is  used  to  prepare  more  volatile  acids  from  their  salts. 
Especially  it  is  used  in  the  preparation  of  hydrochloric  and 
nitric  acids.  In  the  reversible  reactions  represented  by 
the  equations: 

NaCl  +  H2SO4  <=*  NaHSO4  +  HC1 
and 

NaNO3  +  H2SO4  +±  NaHSO4  +  HN03 


SULFATES.     SODIUM  THIOSULFATE  125 

the  fact  that  hydrochloric  acid  is  a  gas  and  nitric  acid  is 
volatile  at  a  moderate  temperature  carries  the  reaction  to 
the  right  as  the  hydrochloric  acid  or  the  nitric  acid  escapes 
and  the  operation  can  be  made  practically  quantitative  in 
that  direction. 

Large  quantities  of  sulfuric  acid  are  used  in  the  treatment 
of  the  mineral  phosphates  which  are  used  as  fertilizers.  The 
treatment  makes  the  phosphates  more  soluble  and  so  mo're 
quickly  available  for  the  growth  of  crops.  Sulfuric  acid 
is  also  used  in  making  nitroglycerine  for  dynamite,  in 
making  cellulose  nitrate  or  "gun  cotton, "  which  is  the  basis 
of  smokeless  powder,  in  purifying  petroleum  for  the  manu- 
facture of  kerosene  and  gasolene,  and  in  many  other  ways. 

Sulfates  and  Acid  Sulfates. — Sulfuric  acid  is  dibasic  and 
it  forms  both  normal  and  acid  salts.  Thus  normal  sodium 
sulfate  is  Na2SO4  and  the  acid  sulfate  is  NaHSO4.  The 
acid  salts  are  strongly  acid  in  reaction  and  as  acids  will 
neutralize  bases.  They  are  both  strong  acids  and  salts. 
The  normal  sulfates  of  strong  bases,  such  as  sodium  sul- 
fate, potassium  sulfate  and  calcium  sulfate,  are  neutral^  in 
reaction. 

Nearly  all  sulfates  and  nearly  all  salts  of  othep  strong 
acids,  that  is,  of  acids  which  are  largely  ionized/  in  dilute 
solutions,  are  soluble  in  water.  The  most  commox^insoluble 
sulfates  are  lead  sulfate,  PbSO4,  barium  sulfate,  BaSO4, 
and  calcium  sulfate,  CaSO4.  The  last  is  sufficiently  soluble, 
however,  to  be  a  troublesome  constituent  in  hard  waters. 
Barium  sulfate,  BaS04,  requires  approximately  400,000 
parts  of  water  for  its  solution  and  because  of  this  extreme 
insolubility  it  is  used  to  detect  the  presence  of  sulfuric 
acid  or  a  sulfate  in  solutions  and  to  determine  the  quantity 
of  the  sulfate  group,  SO4,  in  salts  or  mixtures. 

Sodium  Thiosulfate. — A  solution  of  sodium  sulfite  slowly 
absorbs  oxygen  from  the  air  and  the  salt  is  changed  to 
sodium  sulfate: 

Na2S03  +  O  =  Na2SO4 


126  SULFUR,  SELENIUM  AND  TELLURIUM 

The  solution  of  sodium  sulfite  will  dissolve  sulfur,  which 
adds  itself  very  much  as  the  dxygen  does,  forming  a  thio- 
sulf  ate : 

Na2SO3  +  S  =  Na2S203 

Sodium 
thiosulfate 

The  prefix  ihio  (from  Greek  Oeiov,  sulfur)  is  used  to  indi- 
ca£e  that  in  this  compound  we  have  a  sulfate  in  which  one 
atom  of  sulfur  has  taken  the  place  of  an  atom  of  oxygen. 

Sodium  thiosulfate  crystallizes  with  water,  the  salt  hav- 
ing the  composition  Na2S2O3.5H2O.  It  is  much  used  in 
photography  to  "fix"  pictures  by  dissolving  that  portion 
of  the  silver  chloride  or  bromide  which  has  not  been  affected 
by  expos  are  to  the  light.  The  salt  has  been  known  for 
a  long  time  and  when  it  was  first  discovered  the  acid  from 
which  it  was  derived  was  called  hyposulfurous  acid  and  the 
salt  sodium  hyposulfite  because  the  acid  and  salt  contain 
less  oxygen  than  sulfurous  acid  and  sodium  sulfite,  in  pro- 
portion to  the  sulfur  present  in  the  compounds.  While 
chemists  generally  use  the  other  names,  the  name  sodium 
I  yposulfite  is  commonly  used  by  pharmacists  and  photog- 
raphers. 

Selenium  and  Tellurium. — On  referring  to  the  table 
on  p.  100  it  will  be  seen  that  these  two  elements  correspond 
in  their  atomic  weights  to  bromine  and  iodine  of  the  halogen 
family.  They  form  compounds  similar  to  the  compounds 
of  sulfur,  among  which  may  be  mentioned  the  compounds 
having  the  formulas  H2Se,  H2Te,  Se02,  TeO2,  H2Se03, 
H2Te03,  H2Se04,  H2Te04. 

Selenium  occurs  in  several  allotropic  forms,  as  sulfur 
does.  It  is  more  metallic  in  its  properties  than  sulfur  and 
the  electrical  conductivity  of  a  film  of  the  element  is  very 
sensitive  to  the  action  of  light.  This  property  has  been 
used  in  an  instrument  for  measuring  the  intensity  of  light 
from  very  distant  sources  and  has  been  applied  in  the  study 
of  variable  stars. 


SUMMARY.     SULFUR  127 

Tellurium  is  still  more  metallic  and  in  many  of  its  proper- 
ties it  should  be  classed  with  the  metals.  It  is  sometimes 
found  associated  with  gold,  telluride  gold  ores  being  found 
especially  in  Colorado. 

SUMMARY 

Sulfur  is  found  free  in  nature,  also  combined  in  the  sulfides, 
iron  pyrites,  galena  and  sphalerite  and  the  sulfates,  gypsum 
and  barium  sulfate. 

Sulfur  is  used  in  making  sulfuric  acid,  in  "sulfuring" 
fruit  for  drying,  in  the  "lime-sulfur"  wash,  in  black  gun 
powder,  in  the  manufacture  of  india-rubber  and  in  making 
some  dyes. 

Sulfur  has  three  allotropic  solid  forms,  rhombic  and 
needle-shaped  crystals  and  an  amorphous  form. 

There  are  two  liquid  forms,  corresponding  to  the  last 
two  solid  forms. 

Sulfur  boils  at  445°, 

Hydrogen  sulfide  is  found  in  sulfur  waters  and  in  decayed 
eggs. 

For  purposes  of  anaylsis  metals  are  separated  by  the 
conduct  of  solutions  of  their  salts  toward  hydrogen  sulfide. 
Metals  of  Group  I  are  precipitated  by  hydrogen  sulfide 
from  acid  solutions,  those  of  Group  II  from  alkaline  solu- 
tions, and  those  of  Group  III  are  precipitated  from  neither 
acid  nor  alkaline  solutions. 

The  formation  of  a  precipitate  is  a  reversible  reaction 
which  goes  nearly  to  completion  because  the  precipitate  is 
removed  from  the  sphere  of  action. 

Hydrosulfides  are  formed  by  the1  action  of  hydrogen  sul- 
fide on  hydroxides. 

Sulfur  dioxide  is  prepared  by  burning  sulfur  or  by  the 
action  of  concentrated  sulfuric  acid  on  copper.  It  is  used 
in  bleaching,  as  a  germicide  and  in  the  manufacture  of 
sulfuric  acid. 


128  SULFUR,  SELENIUM  AND  TELLURIUM 

Sulfurous  acid  is  formed  when  sulfur  dioxide  is  dissolved 
in  water.  It  is  a  bibasic  acid,  giving  both  acid  and  normal 
salts. 

Sulfur  trioxide  is  a  liquid  which  easily  polymerizes  to  a 
solid  form.  It  combines  with  water  to  form  sulfuric 
acid. 

In  accordance  with  the  principle  of  van't  Hoff-Le  Chate- 
lier  a  rise  in  temperature  is  unfavorable  to  that  direc- 
tion of  a  reversible  reaction  in  which  heat  is  evolved. 

The  chamber  process  for  sulfuric  acid  is  carried  out  in 
leaden  chambers  into  which  are  introduced  sulfur  dioxide, 
air,  nitric  acid  and  water. 

Nitric  acid  or  nitrogen  dioxide  oxidizes  sulfur  dioxide  and 
water  to  sulfuric  acid  and  the  nitric  oxide  formed  combines 
with  oxygen,  acting  in  this  way  as  a  catalytic  agent  to  con- 
vert sulfur  dioxide  into  sulfuric  acid. 

Sulfuric  acid  is  concentrated  in  glass,  platinum  or  iron. 
It  has  a  strong  affinity  for  water. 

Sulfuric  acid  is  used  in  preparing  hydrochloric  and  nitric 
acids,  in  making  smokeless  powder,  in  purifying  petroleum 
and  in  many  other  ways. 

Sulfuric  acid  is  bibasic.  The  sulfates  of  calcium,  lead 
and  barium  are  only  slightly  soluble. 

Sodium  thiosulfate  is  prepared  by  dissolving  sulfur  in  a 
solution  of  sodium  sulfite.  It  is  used  to  "fix"  pictures  in 
photography. 

Selenium  and  tellurium  are  elements  of  higher  atomic 
weights  having  properties  similar  to  those  of  sulfur. 

EXERCISES 

1.  A  liter  of  hydrogen  sulfide  weighs  approximately  1.52  grams. 
How  much  ferrous  sulfide  and  how  much  sulfuric  acid  will  be 
required  Jbo  furnish  22.4  cc.  of  the  gas? 

2.  How  many  pounds  of  sulfur  will  be  required  to  make  a  ton  of 
sulfuric  acid  of  70  per  cent  strength?     How  much  to  make  a  ton 
of  95  per  cent  acid? 


EXERCISES.     SULFUR 

3.  How  many  pounds  of  iron  pyrites  will  be  required  to  make  a 
ton  of  pure  sulfuric  acid? 

4.  How  much  sulfuric  acid  will  be  required  to  give  a  ton  of  pure 
nitric  acid,  HNOs? 

5.  If  25  pounds  of  sodium  nitrate  are  used  in  making  a  ton  of 
95  per  cent  sulfuric  acid,  how  many  times  does  each  atom  of  nitro- 
gen react  in  the  process? 

6.  Heat  is  evolved  when  chlorine  is  liberated  by  the  action  of 
oxygen  or  air  on  hydrochloric  acid  (p.  82).     Will  an  increase  in 
temperature  favor  the  formation  of  chlorine  or  the  reverse? 

7.  Name  the  compounds  whose  formulas  are  given  in  the  para- 
graph on  selenium  and  tellurium. 


CHAPTER  XII 


SELECTION  OF  MOLECULAR  AND  ATOMIC  WEIGHTS 

Combination  of  Gases  by  Volume. — Two  volumes  of 
hydrogen  unite  with  one  volume  of  oxygen  to  form  water. 
If  the  water  formed  is  kept  at  the  same  temperature  as  the 
oxygen  and  hydrogen  which  combine  and  this  temperature 
is  high  enough  so  that  the  water  remains  as  a  gas,  that  is,  in 
the  form  of  steam,  two  volumes  of  steam  will  result  from 
the  combination.  One  volume  of  hydrogen  combines  with 
one  volume  of  chlorine  to  form  two  volumes  of  hydro- 
chloric acid,  HC1.  One  volume  of  nitrogen  unites  with 
three  volumes  of  hydrogen  to  form  ammonia,  NH3.  One 
volume  of  chlorine  unites  with  two  volumes  of  oxygen  to 
form  two  volumes  of  chlorine  dioxide,  C1O2.  These  rela- 
tions can  be  represented  by  the  following  diagrams: 


2  grams 
Hydrogen 

+ 

71  grams 
Chlorine 

2  grams 
Hydrogen 

32  grams 
Oxygen 

2  grams 
Hydrogen 

36.5  grams 
Hydrochlor- 
ic acid 

36.5  grams 
Hydrochlor- 
ic acid 

18  grams 
Steam 


18  grams 
Steam 


130 


GAY  LUSSAC'S  LAW 


131 


2  grams 
Hydrogen 

+ 

2  grams 
Hydrogen 

28  grams 
Nitrogen 

2  grams 
Hydrogen 

32  grams 
Oxygen 

71  grams 
Chlorine 

32  grams 
Oxygen 

17   grams 
Ammonia 

17  grams 
Ammonia 

67.5  grams 

Chlorine 

dioxide 


67.5  grams 

Chlorine 

dioxide 


Gay  Lussac's  Law. — 4  study  of  the  combination  of  gases 
with  each  other  led  to  the  discovery  of  Gay  Lussac's  Law 
of  Combination  by  Volume.  There  is  always  a  simple 
ratio  between  the  volumes  of  gases  which  combine  and  there  is 
also  a  simple  ratio  between  these  volumes  and  the  volume  of 
the  product  if  that  is  a  gas. 

Relations  between  Combining  Volumes  and  Atomic 
Weights. — The  volumes  for  the  diagrams  have  been  selected 
in  such  a  manner  that  each  volume  of  a  gas  containing 
hydrogen  contains  either  one  gram  or  a  whole  number  of 
grams  of  that  element.  If  the  unit  volume  is  selected  in 
this  way  it  follows: 


132  MOLECULAR  AND  ATOMIC  WEIGHTS 

1.  The  ratio  of  the  volumes  is  always  one  of  simple  whole 
numbers  (Gay  Lussac's  law). 

2.  The  composition  of  every  compound  must  be  expressed 
in   terms  of  the  atomic  weights  of  the  elements  multiplied  by 
small  numbers  (law  of  combining  proportions,  p.  48). 

3.  This  volume  must  either  contain  one  gram  atom  or  a  whole 
number  of  gram  atoms,  or  a  simple  fraction  of  a  gram  atom 
of  every  other  element  present. 

As  hydrogen  is  the  unit1  for  atomic  weights  it  is  convenient 
to  select  as  our  unit  for  the  combining  of  gases  that  volume 
of  some  compound  containing  one  atom  of  hydrogen  in  the 
molecule,  which  contains  one  gram  atom  of  hydrogen.  We 
shall  be  likely  to  discover  such  a  compound  if  we  examine  a 
large  number  of  compounds  of  hydrogen  and  select  the  one 
which  contains  the  smallest  weight  of  hydrogen  in  a  given 
volume.  Among  the  compounds  chosen  for  illustration 
hydrochloric  acid  contains  the  smallest  amount  of  hydrogen 
in  the  unit  volume.  An  examination  of  many  hundreds  of 
other  compounds  of  hydrogen  has  revealed  none  which  con- 
tains less  than  one  gram  of  hydrogen  in  this  volume.  We 
may  feel  reasonably  certain,  therefore,  that  hydrochloric 
acid  has  only  one  atom  of  hydrogen  in  its  molecule  and 
the  volume  of  hydrochloric  acid  which  contains  one  gram 
atom  of  hydrogen  may  be  selected  as  our  unit  volume.  This 
volume  is  22.4  liters  (36.5  divided  by  1.6398,  the  weight 
of  one  liter  of  hydrochloric  acid  gas  at  0°  and  760  mm. 
=  22.3). 2 

Having  selected  22.4  liters  as  our  unit  volume  it  follows, 
as  has  been  shown,  that  this  volume  of  every  other  gaseous 
compound  of  hydrogen  must  contain  one,  two,  three  or 
some  other  whole  number  of  grams  of  hydrogen  and  that 

1  In  round  numbers.     O  =  16  is  the  real  basis  (see  p.  51). 

2  Oxygen,  nitric  oxide,  carbon  dioxide  and  other  gases  which  are  not  so 
easily  condensed  to  a  liquid  as  hydrochloric  acid  give,  in  the  same  manner, 
values  very  close  to  22.4.     These  gases  obey  the  gas  laws  more  closely 
and  are  considered  "perfect"  gases. 


GRAM  MOLECULAR  VOLUME  133 

it  must  also  contain  one,  two,  three  or  some  other  whole 
number  of  gram  atoms  of  each  element  with  which  the 
hydrogen  is  combined.  This  is  illustrated  in  the  diagrams: 
22.4  liters  of  steam1  contain  2  grams  of  hydrogen  and  16 
grams  of  oxygen;  22.4  liters  of  ammonia  contain  3  grams 
of  hydrogen  and  14  grams  of  nitrogen. 

The  law  of  combining  volumes  enables  us  to  extend  the 
principle  to  compounds  which  do  not  contain  hydrogen, 
also.  Thus  22.4  liters  of  chlorine  dioxide  contain  2  gram- 
atoms,  or  71  grams,  of  chlorine  and  1  gram  atom,  or  16 
grams,  of  oxygen. 

As  22.4  liters  has  been  selected  as  the  unit  volume  because 
no  gaseous  compound  of  hydrogen  contains  less  than  one 
gram  of  the  element  in  that  volume,  so  we  may  select  for 
other  elements  the  smallest  weight  of  the  element  found  in 
22.4  liters  of  the  various  gaseous  compounds  of  the  element. 
If  the  compounds  selected  in  this  manner  contain  only  one 
atom  of  the  element  in  the  molecule,  the  weight  of  the 
element  in  the  unit  volume  must  be  one  gram  atom,  and 
the  atomic  weight  of  the  element  follows  from  this.  Thus 
there  is  no  gaseous  compound  of  oxygen  which  contains 
less  than  16  grams  of  the  element  in  22.4  liters.  The  atomic 
weight  of  oxygen  is  selected  as  16,  therefore,  and  in  a 
similar  manner  35.5  has  been  chosen  as  the  atomic  weight 
of  chlorine,  14  as  the  atomic  weight  of  nitrogen  and  32  as 
the  atomic  weight  of  sulfur.  Other  atomic  weights  have 
been  chosen  in  the  same  way. 

The  fact  that  the  atomic  weights  chosen  by  the  principles 
which  have  been  outlined  have  led  to  an  orderly  arrange- 
ment of  the  elements  (the  Periodic  System,  p.  162)  in  which 
there  has  been  shown  to  be  a  close  relation  between  the 
atomic  weights  and  the  properties  of  the  elements  is  a  very 

1  By  this  is  meant,  of  course,  that  volume  of  steam  which  would  fill 
22.4  liters  if  it  could  be  brought  to  normal  conditions  (0°  and  760  mm.) 
without  condensation  to  a  liquid. 


134  MOLECULAR  AND  ATOMIC  WEIGHTS 

strong  reason  for  believing  that  these  principles  are  correct 
and  that  the  numbers  chosen  actually  represent  the  rela- 
tive weights  of  the  atoms  of  the  elements. 

Gram  Molecular  Volume. — The  molecular  wei-ght  of  a 
compound  is  the  sum  of  the  atomic  weights  of  the  elements 
of  which  it  is  composed,  each  atomic  weight  being  taken  as 
many  times  as  there  are  atoms  of  the  element.  Thus  the 
molecular  weight  of  water,  H2O,  is  2  +  16  =  18;  the  mole- 
cular weight  of  ammonia,  NH3,  is  3  +  14  =  17;  the  mole- 
cular weight  of  chlorine  dioxide,  C1O2,  is  35.5  +  (2  X  16) 
=  67.5;  the  molecular  weight  of  sulfuric  acid,  H2SO4, 
is  (2  X  1)  +  32  +  (4  X  16)  =  98.  A  gram  molecule  of  a 
compound  weighs  as  many  grams  as  the  molecular  weight. 
A  gram  molecule  of  water  weighs  18  grams.  A  gram  mole- 
cule of  sulfuric  acid  weighs  98  grams. 

It  has  been  shown  in  the  preceding  paragraph  that  the 
unit  volume,*  22.4  liters,  has  been  so  chosen  that  it  must 
contain  as  manjr  gram  atoms  of  hydrogen  and  of  each  of 
the  other  elements  as  there  are  atoms  of  the  element  in 
the  compound.  Thus  22.4  liters  of  steam  contain  (2X1) 
+  16  =  18  grams  of  steam  and  22.4  liters  of  chlorine  dioxide 
contain  35.5  +  (2  X  16)  =  67.5  grams  of  the  gas.  It 
follows,  then,  that  22.4  liters  of  any  gas  contain  one  gram 
molecule  of  the  gas.  Because  of  this  relation,  22.4  liters  is 
called  a  gram-molecular  volume.  If  we  can  determine  the 
weight  of  22.4  liters  of  any  gas  under  standard  conditions1 
we  have  the  weight  of  a  gram  molecule  and  so  the  molecular 
weight  of  the  compound  or  element.  Conversely,  if  we 
know  the  formula  of  any  gaseous  compound  we  have  only 
to  add  the  weights  of  the  atoms  together  to  find  the 
weight  in  grams  of  22.4  liters  of  the  compound. 

An  additional  meaning  which  may  be  given  to  formulas 

1  At  0°  and  760  mm.  If  the  determination  is  made  at  any  other  tem- 
perature or  pressure,  the  result  may  be  calculated  for  standard  conditions 
by  means  of  the  gas  laws  (p.  37). 


AVOGADRO'S  LAW  135 

of  gaseous  compounds  follows  from  the  foregoing.  Just 
as  a  formula  may  represent  a  gram  molecule  of  a  compound 
(p.  56)  it  may  also  represent  a  gram^molecular  volume. 
Thus  the  equation: 

N2  +  3H2  =  2NH? 

Nitrogen  Hydrogen    Ammonia 

not  only  means  that  one  molecule  of  nitrogen,  N2  (see 
below),  combines  with  three  molecules  of  hydrogen,  H2,  to 
form  two  molecules  of  ammonia,  NH3,  but  it  also  means  that 
a  gram-molecular  volume  (22.4  liters)  of  nitrogen  combines 
with  three  gram-molecular  volumes  (3  X  22.4  liters)  of 
hydrogen  to  form  two  gram-molecular  volumes  (2  X  22.4 
liters)  of  ammonia. 

Number  of  Molecules  in  the  Unit  Volume  of  a  Gas. 
Avogadro's  Law.— If  22.4  liters  of  any  gas  contain  a  gram 
molecule  of  the  gas  it  must  follow  that  if  we  could  decrease 
the  unit  volume  till  it  contained  only  1000  molecules  of  the 
gas  the  same  volume  would  contain  1000  molecules  of  any 
other  gas,  for  the  weights  of  this  small  volume  of  different 
gases  must  still  be  proportional  to  their  molecular  weights. 
In  other  words,  equal  volumes  of  all  gases  under  the  same 
conditions  of  temperature  and  pressure  contain  the  same  num- 
ber of  molecules.  This  is  known  as  Avogadro's  law  because 
it  was  discovered  by  an  Italian  chemist,  Avogadro,  in  1811. 
He  based  his  hypothesis  on  the  law  of  combining  volumes. 
It  has  been  pointed  out  (p.  33)  that  the  diffusion  of  gases 
and  other  phenomena  which  have  led  to  the  kinetic  theory 
of  gases  are  best  explained  on  the  supposition  that  equal 
volumes  contain  equal  numbers  of  molecules.  So  many 
different  phenomena  are  consistently  explained  on  this 
basis  that  the  law  is  now  considered  to  be  thoroughly 
established. 

The  actual  number  of  molecules  in  a  given  molecular 
volume  must  be,  from  the  figures  given  on  p.  33,  22.4  X 
1000  X  2.71  X  1019  (or  607,000,000,00Q>QQQ,OQp,000,000).. 


136  MOLECULAR  AND  ATOMIC  WEIGHTS 

Formulas  of  Oxygen  and  of  Other  Elementary  Gases. — 
The  gram-molecular  volume  of  oxygen  weighs  32  grams. 
From  this  the  formula  of  oxygen  must  be  O2.  In  a  similar 
manner  the  formulas  of  hydrogen,  chlorine  and  nitrogen  are 
H2,  C12  and  N2. 

The  same  conclusions  may  be  reached  as  follows  on  the 
basis  of  Avogadro's  law: 

Suppose  that  a  volume  of  hydrogen  containing  1000  mole- 
cules be  taken.  It  will  combine  with  the  same  volume  of 
chlorine,  which  must,  according  to  the  law,  also  contain 
1000  molecules.  The  combination  will  give  two  volumes  of 
hydrochloric  acid,  which  must  contain  2000  molecules. 
In  other  words,  one  molecule  of  hydrogen  with  one  molecule 
of  chlorine  gives  two  molecules  of  hydrochloric  acid. 

1000  +   1000   =    1000  +  1000 

Molecules      Molecules      Molecules      Molecules 

of  of  of  of 

hydrogen         chlorine    hydrochloric  hydrochloric 
acid  acid 

Since  each  molecule  of  hydrochloric  acid  must  contain 
at  least  one  atom  of  hydrogen  and  one  atom  of  chlorine, 
each  molecule  of  hydrogen  and  each  molecule  of  chlorine 
must  contain  at  least  two  atoms  of  the  element. 


SUMMARY 

The  volumes  of  gases  which  react  always  bear  a  simple 
ratio  to  each  other  and  the  ratios  between  these  volumes 
and  the  volume  of  the  product,  if  that  is  a  gas,  is  also 
simple  (Gay  Lussac's  Law). 

Twenty-two  and  four-tenths  liters  of  any  gaseous 
compound  under  standard  conditions  contain  one  gram 
molecule  of  the  compound  and  this  volume  is  called  a 
gram-molecular  volume. 

The  smallest  number  of  grams  of  an  element  found  in  any 


EXERCISES.     MOLECULAR  WEIGHTS  137 

of  its  gaseous  compounds  is  chosen  as  the  weight  of  a 
gram  atom  of  the  element  and  this  fixes  the  atomic  weight.1 
The  formulas  of  oxygen,  hydrogen,  chlorine  and  nitrogen 
are  O2,  H2,  C12,  N2. 

EXERCISES 

1.  What  is  the  weight  of  a  gram-molecular  volume  of  the 
following   gases:  Chlorine   monoxide,    C120,    hydrobromic    acid, 
hydrogen  sulfide,  sulfur  dioxide,  sulfur  trioxide,  hydrogen  tellu- 
ride? 

2.  A  gram-molecular  volume  of  air  weighs  about  29  grams. 
(Notice  the  relation  of  this  to  the  weight  of  a  gram-molecular 
volume  of  nitrogen,  28  grams,  and  of  oxygen,  32  grams.)     What  is 
the  weight  of  each  of  the  gases  mentioned  as  compared  with  the 
weight  of  the  same  volume  of  air? 

&  The  weight  of  a  liter  of  ozone  is  approximately  2.15  grams; 
what  is  the  formula? 

4.  At  200°  the  weight  of  that  volume  of  iodine  vapor  which 
would  fill  one  liter  if  it  could  be  cooled  to  0°  under  a  pressure  of 
760  mm.  is  approximately  11.3  grams.     What  is  the  formula  of 
iodine  at  that  temperature? 

5.  At  1000°  the  weight  of  that  volume  of  iodine  vapor  which 
would  occupy  one  liter  under  standard  conditions  is  5.7  grams. 
What  is  the  formula  of  iodine  at  that  temperature? 

6.  What  weight  of  sulfur  will  be  required  to  give  22.4  liters  of 
sulfur  dioxide?     What  volume  of  oxygen? 

7.  What   volume   of   oxygen    will    be   required    to   give   22.4 
liters  of  sulfur  trioxide,  if  it  could  exist  as  a  gas  under  standard 
conditions? 

1  Approximately.  Accurate  atomic  weights  are  determined  by  the 
quantitative  analysis  of  compounds  of  the  element,  especially  of  com- 
pounds with  chlorine,  bromine  or  iodine.  A  few  elements  have  no  gaseous 
compounds,  or  have  none  containing  a  single  atom  of  the  element  in  a 
molecule.  The  atomic  weights  of  these  elements  are  fixed  by  the  law  of 
Dulong  and  Petit  (p.  266). 


CHAPTER  XIII 
GROUP  V:  NITROGEN 

Occurrence  of  Nitrogen. — Nitrogen  forms  nearly  four- 
iifths  of  the  volume  of  air  and  a  little  more  than  three- 
fourths  of  its  weight.  As  the  pressure  of  the  air  at  sea 
level  is  nearly  15  pounds  to  the  square  inch,  the  weight  of 
nitrogen  above  one  square  inch  is  about  11  pounds,  or  more 
than  three-fourths  of  a  ton  over  one  square  foot  and  more 
than  20,000,000  tons  over  one  square  mile.  In  spite  of  this 
enormous  quantity  of  the  element  in  the  atmosphere  it  is  so 
difficult  to  cause  nitrogen  to  combine  with  other  elements 
that  compounds  of  nitrogen  are  comparatively  expensive. 

Nitrogen  is  an  essential  constituent  of  all  living  bodies, 
whether  animal  or  vegetable.  Corn,  wheat,  oats,  grass  and 
most  vegetables  must  secure  nitrogen  for  their  growth 
from  compounds  of  nitrogen  present  in  the  soil,  especially 
from  potassium  nitrate,  KNO3,  calcium  nitrate,  Ca(NO3)2, 
or  ammonia,  NH3.  Apparently  none  of  the  cereals  can 
utilize  nitrogen  of  the  air  directly  and  they  will  not  grow 
on  a  soil  that  does  not  contain  compounds  of  nitrogen  avail- 
able for  their  use. 

Alfalfa,  clover  and  some  other  legumes  may,  however, 
utilize  the  nitrogen  of  the  air  with  the  aid  of  bacteria  which 
grow  on  their  roots,  producing  nodules.  A  small  amount  of 
-oxygen  and  nitrogen  combine  during  every  lightning  flash 
and  the  oxides  of  nitrogen  formed  combine  with  water  and 
.ammonia  in  the  air  to  form  ammonium  nitrite,  NH4NO2. 
This  is  carried  down  by  rain  and  becomes  available  for  the 
.growth  of  plants. 

Organic  matter  in  the  soil  decays  more  or  less  rapidly 

138 


NATURAL  HISTORY  OF  NITROGEN  1391 

under  the  influence  of  bacteria.  If  air  is  present,  potassium 
nitrate,  KNO3,  or  calcium  nitrate,  Ca(N03)2,  is  formed. 
In  the  absence  of  air,  where  the  conditions  favor  reduction 
rather  than  oxidation,  ammonia,  NH3,  is  formed.  Either 
of  these  may  then  be  available  for  the  growth  of  new  mate- 
rial. The  course  of  nitrogen  in  nature  which  has  been  out- 
lined may  be  seen  from  the  following  diagram: 

Leguminous  plants  with 


the  help  of  bacteria 
Atmospheric 
electricity  Plants 


Atmospheric 
Nitrogen 


Nitrates 


Organic  compounds 
of  Nitrogen 


Denitrifying      NiTtrifying   '  Decayjof  plant" 

bac  teria  and  ani  mal  tissues 

or  dis  I  tillation 

• —        — |  Nitrifying 

Nitrites  U 

bacteria 


Preparation  and  Properties  of  Nitrogen. — When  phos- 
phorus is  burned  in  a  confined  portion  of  air  the  oxygen  is 
removed  and  nearly  pure  nitrogen  remains.  The  oxygen 
may  also  be  completely  removed  from  air  by  passing  it 
over  heated  metallic  copper.  Nitrogen  is  colorless  and 
odorless  and  will  not  support  combustion.  It  is  very 
much  more  inert  than  the  elementary  gases  and  other 
elements  which  have  been  studied  and  this  is  evidently 
one  reason  why  such  large  quantities  of  nitrogen  are  found 
uncombined  in  the  atmosphere. 

At  the  very  high  temperatures  produced  by  electric 
sparks  a  portion  of  the  nitrogen  and  oxygen  of  the  air 
through  which  the  sparks  are  passed  combines  to  form 
nitric  oxide,  NO: 

N2  +  O2  =  2NO 
Nitrogen  and  hydrogen  will  combine  at  moderate  tem- 


140  NITROGEN 

peratures  with  the  aid  of  metallic  osmium  or  uranium  as  a 
catalyzer : 

N2  +  3H2  =  2NH3 

Both  of  these  processes  are  now  in  technical  use  and 
promise  to  be  increasingly  important. 

Ammonia. — When  organic  matter,  such  as  stable  manure, 
or  sewage,  decomposes  with  exclusion  of  air,  through  the 
action  of  bacteria,  a  part  of  the  nitrogen  is  converted  into 
ammonia,  NH3.  Ammonia  is  also  formed  when  almost  any 
kind  of  natural  organic  matter  containing  nitrogen  is  heated 
to  a  high  temperature.  In  this  manner  ammonia  is  formed 
when  coal  is  heated  for  the  manufacture  of  illuminating  gas 
or  of  coke  and  nearly  all  of  the  ammonia  and  ammonium 
salts  of  commerce  come  from  this  source. 

Ammonia  is  most  readily  prepared  for  laboratory  pur- 
poses by  warming  a  concentrated  solution  of  the  gas  in 
water.  The  gas  may  be  dried  by  passing  it  through  a  tube 
filled  with  solid  potassium  hydroxide  or  sodium  hydroxide, 
or  with  quicklime. 

Ammonia  is  a  colorless  gas  with  a  pungent  odor.  Is  the 
gas  heavier  or  lighter  than  air,  and  in  what  proportion? 
A  gram-molecular  volume  of  air  weighs  about  29  grams. 

Water  at  the  freezing  point  will  dissolve  1000  times  its 
volume  of  ammonia.  The  solution  is  lighter  than  water 
and  gives  off  ammonia  very  readily  when  it  is  warmed. 
The  solution  of  ammonia  reacts  alkaline  toward  litmus  and 
neutralizes  acids.  These  properties  and  others  have  led 
to  the  view  that  the  ammonia  combines  with  the  water,  in 
part,  to  form  ammonium  hydroxide,  NH4OH : 

NH3  +  HOH  =  NH4OH 
or,  graphically, 

H\  H\     /^ 

H^N  +  H— O— H   =   H-^N(^ 

W  W       X0— H 


AMMONIUM  SALTS  14 1 

The  nitrogen,  which  is  trivalent  in  ammonia,  becomes 
quinquivalent  in  ammonium  hydroxide.  The  ammonium 
hydroxide  gives  in  a  solution  the  positive  ammonium  ion, 
NH4+,  and  the  negative  hydroxide  ion,  OH~.  It  is  the 
hydroxide  ion,  of  course,  which  causes  the  alkaline  reaction 
of  the  solution  (p.  74). 

Ammonium  Salts. — Ammonia  combines  directly  with 
acids  to  form  ammonium  salts  in  which  the  univalent 
ammonium  group,  NH4,  takes  the  place  usually  occupied 
by  a  metal: 

NH3  +  HC1  =  NH4C1 

Ammonium  chloride 

Ammonium  chloride  resembles  common  salt,  NaCl,  in 
taste  and  in  many  of  its  properties.  How  would  it  react 
with  concentrated  sulfuric  acid? 

2NH3  +  H2S04  =  (NH4)2SO4 

Ammonium  sulfate 

NH3  +  HN03  =  NH4N03 

Ammonium  nitrate 

Or,  graphically, 

H\  H\     /H 

H^N  +  H-C1      =   H^N( 

W  W       XC1 

In  each  case  one  molecule  of  ammonia  combines  with  one 
atom  of  hydrogen  from  the  acid.  For  the  formation  of  a 
normal  salt  a  monobasic  acid  requires  one  molecule  of 
ammonia,  a  bibasic  acid  two  molecules  and  a  tribasic  acid 
three  molecules. 

Action  of  Bases  on  Ammonium  Salts. — When  an  ammo- 
nium salt  is  treated  with  a  base  the  usual  interaction  with 
partial  exchanges  of  ions  takes  place: 

NH4C1  +  NaOH  <=»  NaCl  +  NH4OH 
If   little   water   is   present,    the    ammonium   hydroxide 
decomposes  and  the  ammonia  escapes: 
NH4OH  *±  NH3  +  HO 


142  NITROGEN 

This  is  the  usual  commercial  method  of  preparing 
ammonia,  but  calcium  hydroxide,  Ca(OH)2,  is  used  instead 
of  sodium  hydroxide.  Why? 

Synthesis  of  Ammonia. — It  has  been  known  for  a  long 
time  that  when  electric  sparks  are  passed  through  a  mixture 
of  nitrogen  and  hydrogen  a  minute  quantity  of  ammonia 
is  formed: 

N2  +  3H2  <=±  2NH3 

But  ammonia  gas  is  also  decomposed  into  nitrogen  and 
hydrogen  by  the  passage  of  electric  sparks  and  the  equilib- 
rium in  this  reversible  reaction  is  so  very  far  on  the  side 
toward  decomposition  that  it  is  wholly  impracticable  to 
prepare  ammonia  by  this  method. 

A  practical  commercial  synthesis  was  finally  discovered 
by  a  careful  study  and  application  of  the  following  facts 
and  principles: 

First.- — The  reaction  is  exothermic,  that  is,  heat  is  gen- 
erated when  nitrogen  and  hydrogen  combine. 

Second. — Because  heat  is  generated  in  the  combination, 
an  increase  in  the  temperature  favors  the  decomposition 
&iid  not  the  formation  of  ammonia  (principle  of  Van't 
Hoff-Le  Chatelier,  p.  120). 

Third. — The  combination  of  the  gases  takes  place  very 
slowly  at  any  temperature  at  which  the  equilibrium  would 
be  favorable;  that  is,  at  any  temperature  at  which  the 
combination  goes  far  enough  to  be  commercially  profitable. 

Fourth. — Osmium,  uranium  and  some  other  metals 
catalyse  the  reaction  (p.  9)  and  make  it  practicable  at  a 
temperature  of  400°  or  below. 

Fifth. — Four  volumes  of  the  mixture  of  nitrogen  and 
hydrogen  give  only  two  volumes  of  ammonia.  An  increase 
of  the  pressure,  therefore,  favors  the  combination  (prin- 
ciple of  Van't  Hoff-Le  Chatelier). 

On  the  basis  of  .these  facts  and  principles  the  commercial 


OXIDATION  OF  AMMONIA  143 

manufacture  of  ammonia  is  carried  out  by  passing  a  highly 
compressed  mixture  of  nitrogen  and  hydrogen  over  osmium, 
uranium  or  some  other  catalyzer.  A  commission  of  the 
Government  of  the  United  States  has  recommended  the 
use  of  this  process  on  a  large  scale  for  the  manufacture  of 
ammonia,  which  is  then  to  be  oxidized  to  nitric  acid  (see 
below).  There  is  strong  reason  for  believing  that  this  and 
other  synthetic  processes  for  making  ammonia  and  nitric 
acid  were  extensively  used  in  Germany  during  the  Great 
War  and  that  it  is  pnly  through  the  use  of  these  processes 
that  her  speedy  defeat,  through  lack  of  munitions,  was 
prevented. 

Oxidation  of  Ammonia  to  Nitric  Oxide  and  Nitric  Acid. — 
Smokeless  powder,  gun  cotton,  nitroglycerine,  trinitro- 
toluene ("T.  N.  T.  ")  gunpowder  and  all  other  explosives, 
used  in  blasting  or  in  warfare,  except  some  of  the  primers, 
and  "ammonal,"  a  mixture  of  ammonium  chloride  and 
aluminum  powder,  require  nitric  acid  or  a  nitrate  for  their 
manufacture.  Before  the  nineteenth  century  the  world's 
supply  of  nitric  acid  and  nitrates  was  obtained  from  salt- 
peter, KNO3,  and  a  few  other  natural  nitrates.  For  many 
years  before  the  Great  War  the  principal  source  of  supply 
was  Chili  saltpeter,  NaNOa.  This  natural  supply  will  be 
exhausted  within  a  comparatively  few  years  and  chemists 
have  been  searching  eagerly  for  methods  of  producing  nitric 
acid  from  other  materials.  One  of  the  most  promising  of 
the  methods  discovered  was  developed  by  Ostwald.  It 
consists  in  passing  a  mixture  of  ammonia  and  air  over 
heated  platinized  asbestos,  which  serves  as  a  catalyzer. 
The  hydrogen  of  the  ammonia  is  burned  to  water  and  the 
nitrogen  to  nitric  oxide.  The  latter  can  be  converted  to 
nitric  acid  without  much  difficulty  (p.  151) : 
4NH3  +  5O2  =  6H20  +  4NO 

About  90  per  cent  of  the  nitrogen  of  the  ammonia  may  be 
converted  into  nitric  acid  by  this  process. 


144  NITROGEN 

Nitric  Acid. — When  organic  matter  decays  under  the 
influence  of  bacteria  in  a  soil  to  which  air  has  free  access 
a  part  of  the  nitrogen  is  converted  into  nitrates.  The 
nitrates  most  common  in  the  soil  are  those  of  potassium, 
sodium  and  calcium,  KNOs,  NaNO3  and  Ca(NO3)2. 

In  an  arid  region  in  Chili  and  Peru,  in  South  America, 
enormous  beds  of  sodium  nitrate,  NaNO3,  have  been  formed 
and  for  many  years  this  mineral  has  been  the  almost  ex- 
clusive source  from  which  nitric  acid  and  the  saltpeter  or 
potassium  nitrate,  KNO3,  used  for  gunpowder,  have  been 
manufactured  (see  above). 

Nitric  acid  is  manufactured  by  mixing  sodium  nitrate 
with  concentrated  sulf uric  acid  and  distilling  the  mixture : 

NaN03  +  H2S04  =  NaHSO4  +  HNO3 

Acid  sodium 
sulfate 

Nitric  acid  is  volatile  at  a  much  lower  temperature  than 
sulf  uric  acid  and,  because  of  this,  the  reversible  reaction 
may  be  carried  to  completion  in  the  direction  which  is 
desired. 

Properties  of  Nitric  Acid. — Pure  nitric  acid  boils  at  86°. 
This  pure  acid  is  difficult  to  prepare  and  decomposes  easily, 
especially  when  exposed  to  the  light: 

4HN03  =  4NO2  +  O2  +  2H2O 

Nitrogen 
dioxide 

The  nitrogen  dioxide  gives  a  yellow  color  to  nitric  acid 
containing  it. 

The  ordinary  concentrated  nitric  acid  of  the  laboratory 
contains  30  to  35  per  cent  of  water.  Such  an  acid  boils 
at  about  120°  and  is  very  much  more  stable  than  the  more 
concentrated  acid. 

Oxidation  with  Nitric  Acid. — Pure  nitric  acid  is  a  very 
powerful  oxidizing  agent.  If  a  little  of  the  acid  is  put  in  a 
test-tube  and  some  wool,  hair  or  feathers  is  placed  in  the 
mouth  of  the  tube,  on  boiling  the  acid  the  wool,  or  the  other 


NITRIC  ACID  AND  METALS  145 

materials  mentioned,   will  take  fire  and  burn  when  the 
vapor  of  the  acid  comes  in  contact  with  them. 

Burning  charcoal  will  continue  to  burn  if  thrust  beneath 
the  surface  of  the  pure  acid. 

Tin  is  oxidized  by  ordinary  strong  nitric  acid  to  meta- 
stannic  acid,  a  white  powder  containing  hydrogen  and  oxy- 
gen as  well  as  tin.  The  nitric  acid  is  reduced  to  nitric  oxide, 
NO,  and  nitrogen  dioxide,  N(>2. 

Action  of  Nitric  Acid  on  Metals. — From  the  action  of 
hydrochloric  or  sulfuric  acid  on  zinc  and  iron  we  should 
expect  that  nitric  acid  would  give,  with  these  metals,  nitrates 
and  hydrogen.  The  action  of  sulfuric  acid  on  copper, 
however,  has  shown  us  that  the  action  of  an  acid  on  a  metal 
may  take  a  different  course.  In  that  case  the  sulfuric 
acid  is  reduced  and  no  hydrogen  is  liberated  but  sulfur 
dioxide,  a  reduction  product  of  sulfuric  acid,  instead.  A 
high  temperature  is  required  for  sulfuric  acid  to  act  rapidly. 
Nitric  acid  acts  in  a  similar  manner,  at  ordinary  tempera- 
tures, on  all  metals  which  are  attacked  by  it.  This  effect 
is  easily  understood  when  we  remember  that  nitric  acid  is 
much  more  vigorous  than  sulfuric  acid,  as  an  oxidizing 
agent. 

We  may  explain  the  action  of  the  acid  by  supposing  that 
the  nitric  acid  oxidizes  the  metal  and  is  at  the  same  time 
reduced  to  one  of  the  lower  oxides  of  nitrogen  and  water, 
and  that  the  oxide  of  the  metal  reacts  with  more  of  the  acid 
to  form  a  nitrate. 

Or  we  may  suppose  that  the  metal  displaces  the  hydrogen 
of  the  acid  forming  a  nitrate  and  that  the  hydrogen  at  the 
moment  of  liberation  acts  upon  more  of  the  acid,  reducing 
it.  The  following  equations  for  the  action  of  nitric  acid  on 
copper  illustrate  the  two  methods  of  explaining  the  action. 
The  formulas  in  brackets  indicate  substances  which  may 
be  supposed  as  intermediate  in  the  reaction  but  for  whose 

formation  we  have  no  direct  evidence. 
10 


146  NITROGEN 

First  explanation : 

2HNO3  +  3Cu  =  [3CuO]  +  2NO  +  H2O 
6HN03  +  [3CuQ]  =  3Cu(NO3)2  +  3H2O 
Adding  after  eliminating  [3CuO]: 

8HNO3  +  3Cu  =  3Cu(NO3)2  +  2NO  +  4H20 
Second  explanation: 

2HNO3  +  Cu     =  Cu(NO3)2  +  [2H] 
2HNO3  +  [6H]  =  2NO  +  4H2O 

Multiplying  the  first  equation  by  3  and  adding  it  to  the 
second  after  eliminating  [6H]  we  have : 

8HN03  +  3Cu  =  3Cu(NO3)2  +  2NO  +  4H2O 

It  will  be  noticed  that  the  final  equation  is  the  same  by 
one  explanation  as  by  the  other  and  that  the  products  of 
the  reaction  give  us  no  means  of  deciding  which  explanation 
is  true.  When  nitric  acid  acts  on  tin  it  oxidizes  the  tin  to 
metastannic  acid  and  tin  nitrate  is  not  formed.  On  the 
other  hand,  when  nitric  acid  acts  on  iron  the  nitric  acid  is 
reduced  to  ammonia,  NH3.  It  seems  that  the  first  expla- 
nation agrees  with  the  action  on  tin  better  than  the  second, 
while  the  second  gives  a  better  account  of  the  action  on  iron. 

Write  the  equations  for  the  following  reactions:  iron 
and  sulfuric  acid;  hydrogen  and  nitric  acid  giving  ammonia; 
ammonia  and  sulfuric  acid;  ferrous  sulfate  and  sodium 
hydroxide  giving  ferrous  hydroxide,  Fe(OH)2;  ammonium 
sulfate  and  sodium  hydroxide. 

Aqua  Regia. — Neither  nitric  nor  hydrochloric  acid  alone 
will  dissolve  gold  or  platinum  but  the  metals  dissolve 
easily  in  a  mixture  of  the  two  acids.  The  nitric  acid  as  an 
oxidizing  agent  liberates  chlorine  from  the  hydrochloric 
acid  and  the  chlorine  combines  with  the  metals  forming 
soluble  chlorides.  The  mixture  of  acids  is  called  aqua 
regia,  meaning  "royal  water."  The  name  refers  to  the 
designation  of  gold  and  platinum  as  " noble"  metals. 


NITROUS  ACID.     NITROUS  OXIDE  147 

Nitrites.  Nitrous  Acid. — When  sodium  nitrate,  NaNO3, 
or  potassium  nitrate,  KN03,  is  heated  with  metallic  lead 
or  copper  either  salt  is  reduced  to  a  nitrite,  NaNO2  or  KNO2. 
The  addition  of  an  acid  to  a  solution  of  a  nitrite  liberates 
nitrous  acid,  HNO2.  If  the  solution  is  dilute,  the  nitrous 
acid  remains  dissolved  in  the  water.  If  the  solution  is 
concentrated,  the  nitrous  acid  decomposes  into  nitrous  an- 
hydride, N203,  and  water.  The  nitrous  anhydride,  in  turn, 
decomposes  into  nitric  oxide,  NO,  and  nitrogen  dioxide, 
NO2. 

Oxides  of  Nitrogen. — There  are  six  oxides  of  nitrogen 
but  two  of  them  have  the  same  composition  and  pass  so 
readily  each  into  the  other  that  they  are  often  spoken  of  as 
a  single  oxide.  The  oxides  are: 

Nitrous  oxide,  N2O 

Nitric  oxide,  NO 

Nitrous  anhydride  N203 , 
f  Nitrogen  dioxide,   NO2 
I  Nitrogen  tetroxide,  N2O4 

Nitric  anhydride,  N2O5 

All  of  these  oxides  except  the  last  may  be  prepared  by 
the  reduction  of  nitric  acid. 

Nitrous  Oxide. — When  ammonium  nitrate,  NH4NO3,  is 
heated  the  hydrogen  of  the  ammonium  group,  NH4,  com- 
bines with  part  of  the  oxygen  of  the  nitrate  group,  NO3, 
while  the  oxygen  remains  combined  with  the  two  nitrogen 
atoms  as  nitrous  oxide,  N20. 

NH4NO3  =  N2O  +  2H2O 

Nitrous  oxide  is  a  colorless  gas  with  a  slightly  sweetish 
taste.  When  inhaled  in  small  quantity  it  produces  an 
intoxicating  effect  and  it  is  called,  for  this  reason,  laughing 
gas.  In  larger  quantities  it  produces  insensibility  and  is 
used  for  minor  surgical  operations,  especially  for  the  extrac- 
tion of  teeth.  By  use  with  oxygen,  to  prevent  suffocation, 


148  NITROGEN 

the  anesthesia  may  be  continued  for  a  longer  time  and  it 
is  claimed  that  the  gas  has  some  marked  advantages  over 
ether  or  chloroform,  which  are  more  often  used. 

Nitrous  oxide  supports  combustion  and  causes  a  glowing 
splinter  to  inflame,  but  its  effect  in  this  regard  is  much  less 
vigorous  than  that  of  oxygen. 

Nitric  Oxide. — The  preparation  of  nitric  oxide  by  the 
action  of  nitric  acid  on  copper  has  been  discussed  and  need 
not  be  described  again  (p.  145). 

At  high  temperatures  nitrogen  and  oxygen  unite  in 
accordance  with  the  reversible  reaction  represented  by  the 
equation : 

N2  +  02  <=±  2NO 

The  reaction  is  endothermic,  that  is,  heat  is  absorbed  as 
the  reaction  proceeds.  It  has  been  pointed  out  that  in 
reversible  reactions  when  heat  is  generated  by  the  combina- 
tion of  elements  or  compounds  an  increase  in  temperature 
shifts  the  equilibrium  toward  the  decomposition  of  the 
product  formed  (principle  of  van't  Hoff-Le  Chatelier, 
pp.  120  and  142).  The  same  principle  leads  to  the  con- 
clusion that  when  heat  is  absorbed  as  the  elements  or  sub- 
stances unite  an  increase  in  the  temperature  causes  the 
equilibrium  to  shift  toward  the  formation  of  the  compound. 
This  conclusion  has  been  confirmed  for  the  combination 
of  oxygen  and  nitrogen  as  shown  in  the  following  table, 
which  gives  the  per  cent  of  the  oxygen  and  nitrogen  com- 
bining when  a  mixture  of  equal  volumes  of  the  gases  is 
heated : 

Temperature  Per  cent  of  NO  calculated          Observed 
1538°  0.35  0.37 

1922°  0.98  0.98 

2402°  2.37  2.23 

2927°  4.43  About  5.00 


NITRIC  OXIDE  149 

As  the  combination  of  nitrogen  and  oxygen  to  form  nitric 
oxide  is  the  first  and  most  difficult  step  in  the  manufacture 
of  nitric  acid  from  atmospheric  nitrogen,  the  conditions 
best  suited  for  the  reaction  have  been  carefully  studied. 
It  is  evident  from  a  consideration  of  the  table  that  the  best 
arrangement  will  be  one  in  which  the  gases  are  brought  to 
the  highest  possible  temperature  and  then  cooled  quickly. 
If  cooled  slowly  the  high  percentage  combination  secured 
at  the  high  temperatures  would  be  lost  because  the  equili- 
brium shifts  toward  the  decomposition  of  the  nitric  oxide 
into  nitrogen  and  oxygen  and  at  these  high  temperatures 
the  reaction  still  goes  rapidly  in  either  direction.  If  the  gas 
can  be  brought  to  a  lower  temperature  without  decomposi- 
tion, the  decomposition  becomes  so  slow  that  the  gas  formed 
at  the  higher  temperatures  is  practically  all  saved.  Even 
at  a  temperature  of  725°  it  is  estimated  that  it  would  take 
80  years  for  one-half  of  the  gas  corresponding  to  the  equili- 
brium to  be  formed  or  decomposed,  while  at  1825°  one-half 
of  the  equilibrium  amount  would  be  formed  or  decomposed 
in  5  seconds — an  illustration  of  the  enormous  acceleration 
of  reactions  at  high  temperatures.  Nitric  oxide  contains 
the  same  amount  of  oxygen  in  a  given  volume  that  nitrous 
oxide,  N2O,  does,  but  the  oxygen  seems  to  be  held  in  a  very 
different  manner.  Nitrous  oxide  causes  a  glowing  splinter 
to  inflame  while  nitric  oxide  will  extinguish  it.  Phosphorus, 
if  well  ignited,  will,  however,  burn  brilliantly  in  nitric 
oxide  and  a  mixture  of  the  vapor  of  carbon  disulfide  with 
nitric  oxide  will  burn  with  a  brilliant  blue  flash. 

Nitric  oxide  is  a  colorless  gas.  It  combines  directly  with 
oxygen  to  form  nitrogen  dioxide,  NC>2,  a  dark  brown  gas 
(see  below). 

Manufacture  of  Nitric  Oxide  and  Nitric  Acid  by  Means  of 
the  Electric  Arc. — For  the  manufacture  of  nitric  oxide  and 
nitric  acid  the  high  temperature  necessary  to  secure  the 
rapid  combination  of  nitrogen  and  oxygen  is  secured  by 


150  NITROGEN 

passing  air  quickly  through  a  long,  flaming,  electric  arc. 
To  secure  the  rapid  cooling  of  the  nitric  oxide,  which  is 
very  essential  in  accordance  with  the  preceding  paragraph, 
the  air,  after  being  heated  in  the  electric  flame,  is  carried 
rapidly  out  of  the  flame  by  the  action  of  a  powerful  magnetic 
field.  In  Norway,  where  water  power  is  cheap  on  account 
of  the  numerous  waterfalls,  it  has  been  found  practicable  to 
manufacture  nitric  oxide  and,  from  this,  nitric  acid,  by 
this  method.  The  consumption  of  energy  is  rather  large 
in  proportion  to  the  quantity  of  nitric  acid  produced  and  it 
is  uncertain  whether  this  process  can  compete  permanently 
with  the  manufacture  of  nitric  acid  by  the  oxidation  of  syn- 
thetic ammonia  (p.  143). 

The  formation  of  nitric  acid  is  completed  by  the  action  of 
the  oxygen  still  remaining  in  the  mixture  and  of  water  (see 
below  under  nitrogen  dioxide) . 

Nitrous  Anhydride. — The  decomposition  of  nitrous  acid, 
HNO2,  into  nitrous  anhydride,  N2O3,  and  the  further  de- 
composition of  nitrous  anhydride  into  nitric  oxide,  NO, 
and  nitrogen  dioxide,  N02,  have  been  referred  to  under 
nitrites  and  nitrous  acid.  Nitrous  anhydride  is  also  formed 
as  a  dark  green  or  blue  liquid  when  nitric  oxide  and  oxygen 
are  brought  together  in  the  right  proportions  at  a  tempera- 
ture below  0°.  In  what  proportion,  by  volume,  should 
nitric  oxide  and  oxygen  be  mixed  to  give  the  compound? 

A  decomposition  like  this,  in  which  a  substance  decom- 
poses and  the  products  of  the  decomposition  recombine 
when  the  conditions  are  reversed,  is  called  dissociation. 
•  Nitrogen  Dioxide,  NO2,  and  Nitrogen  Tetroxide,  N2O4,— 
When  nitric  oxide,  NO,  is  mixed  with  oxygen,  the  two  gases 
combine  directly  to  form  nitrogen  dioxide,  NO2.  In  what 
proportion  should  the  gases  be  mixed?  At  ordinary  tem- 
peratures nitrpgen  dioxide  combines  with  itself,  in  part,  to 
form  the  polymer,  nitrogen  tetroxide,  N2O4.  At  150°  or 
above  this  decomposes  entirely  into  the  dioxide.  How  may 


NITROGEN  TETROXIDE  151 

the  true  formula  at  a  given  temperature  be  determined? 
At  low  temperatures  the  mixture  of  the  dioxide  and 
tetroxide  condenses  to  a  reddish  brown  liquid  and  at 
—  10.5°  this  freezes  to  colorless  crystals,  which  seem  to 
consist  entirely  of  the  tetroxide,  N204. 

Nitrogen  Tetroxide  or  Nitrogen  Dioxide  and  Water. — 
As  nitrous  anhydride,  N2Oa,  dissolves  in  water  forming 
nitrous  acid,  HNO2,  nitrogen  tetroxide  dissolves  in  cold 
water  with  the  formation  of  a  mixture  of  nitric  acid  and 
nitrous  acid.  What  takes  place  in  the  two  cases  is,  perhaps, 
most  easily  understood  by  means  of  the  following  graphical 
formulas : 

Nitrous  anhydride,  0  =  N—  O4-N  =  0 

Water,  H+O  -  H 

or,  N2O3  +  H2O  =  2HN02 

,fl 

Nitrogen  tetroxide,  O  =  N  —  OrN3 


Water,  H-f-OH 

or,  N2O4  +  H2O  =  HNO2  +  HNO3 

With  warm  water  nitrogen  dioxide  gives  nitric  oxide,  NO, 
<*nd  nitric  acid.     This  may  take  place  as  follows: 

Nitrogen  tetroxide,       O  =  N  j  0  — 

: 

Water,  H  -  O   H 


Nitrogen  dioxide, 


. 


The  designation  anhydride  means  "without  water"  and  it 
is  applied  either  to  compounds  which  are  formed  by  the 
removal  of  water  from  an  acid  or  to  compounds  which  com- 
bine with  water  to  form  acids. 


152  NITROGEN 

SUMMARY 

Nitrogen  is  found  free,  in  organic  compounds  and  in 
nitrites  and  nitrates. 

Legumes  with  the  aid  of  bacteria  can  utilize  the  nitrogen 
of  the  air. 

Decomposition  of  organic  matter  by  bacteria  in  the 
absence  of  air  gives  ammonia.  In  the  presence  of  air,  as 
in  a  porous  soil,  it  gives  nitrites  and  nitrates. 

Decomposition  of  organic  matter  by  heat  gives  ammonia. 

Ammonia  is  obtained  as  a  by-product  in  heating  coal  for 
the  manufacture  of  gas  or  coke. 

Ammonia  combines  with  water  to  form  ammonium  hy- 
droxide. It  combines  with  acids  to  form  ammonium  salts. 
Bases  liberate  ammonia  from  such  salts. 

Ammonia  may  be  prepared  by  the  direct  union  of  nitro- 
gen and  hydrogen.  The  reaction  is  exothermic  and  the 
formation  is  favored  by  a  low  temperature  and  by  high 
pressure.  A  catalyzer  must  be  used. 

Ammonia  may  be  oxidized  by  the  oxygen  of  air  to  nitric 
oxide,  with  platinized  asbestos  as  a  catalyzer.  Nitric  acid 
is  used  in  the  manufacture  of  practically  all  explosives. 

Until  recently  nitric  acid  and  nitrates  have  been  prepared 
almost  exclusively  from  Chili  saltpeter. 

Nitric  acid  is  prepared  by  distilling  a  mixture  of  sodium 
nitrate  and  sulfuric  acid. 

Nitric  acid  is  a  powerful  oxidizing  agent.  It  gives  with 
metals  a  nitrate  and  nitric  oxide,  nitrogen  dioxide  or  ammo- 
nia; rarely,  nitrous  oxide.  The  action  may  be  explained  as 
an  oxidation  of  the  metal  or  as  a  reduction  of  the  acid  by 
liberated  hydrogen. 

Aqua  regia  dissolves  gold  or  platinum. 

Sodium  or  potassium  nitrate  may  be  reduced  to  a  nitrite 
by  heating  with  lead  or  copper. 

Nitrites  give  nitrous  acid  with  an  acid.     Nitrous  acid 


EXERCISES.     NITROGEN  153 

decomposes  to  water  and  nitrous  anhydride  and  the  last 
to  nitric  oxide  and  nitrogen  dioxide. 

There  are  six  oxides  of  nitrogen.  All  can  be  prepared 
from  nitric  acid. 

Nitrous  oxide  is  formed  by  heating  ammonium  nitrate. 
It  is  used  as  an  anesthetic. 

Nitric  oxide  is  prepared  by  the  action  of  nitric  acid  on 
copper. 

Nitric  oxide  is  manufactured  by  passing  air  through  an 
electric  arc.  The  reaction  is  endothermic  and  is  favored 
by  a  very  high  temperature  and  rapid  cooling  of  the  product. 

Nitrous  anhydride  is  formed  by  the  combination  of 
nitric  oxide  and  oxygen  at  a  low  temperature.  It  disso- 
ciates to  nitric  oxide  and  nitrogen  dioxide. 

Dissociation  is  a  reversible  decomposition. 

Nitrogen  dioxide  is  formed  by  the  combination  of  nitric 
oxide  and  oxygen. 

Nitrogen  dioxide  gives  nitrous  and  nitric  acids  with 
cold  water;  nitric  acid  and  nitric  oxide  with  hot  water. 

EXERCISES 

1.  What  are  the  acids  which  may  be  formed  from  the  following 
anhydrides:  C120,    C1207,    I205,    S02,    SO3,    P205.      Phosphoric 
anhydride,t  P205,  forms  three  different  acids  by  the  addition  of 
one,  two  or  three  molecules  of  water.     Write  the  equations  for  the 
formation  of  each  of  these. 

2.  What  is  the  weight  of  22.4  liters  of  the  mixture  of  nitric  oxide 
and  nitrogen  dioxide  which  would  combine  to  form  nitrous  an- 
hydride? 

3.  How  many  grams  of  copper  will  be  required  to  prepare  a 
gram-molecular  volume  of  nitric  oxide?     How  many  grams  of 
nitric  acid,  of  sp.  gr.  1.20  and  containing  32  per  cent  of  the  pure 
acid,  will  be  required? 

4.  How  many  grams  of  copper  will  be  required  to  reduce  10 
grams  of  sodium  nitrate  to  sodium  nitrite?     How  many  grams  of 
lead  would  be  required  for  the  same  purpose?     At  the  current 


154  NITROGEN 

market  price  which  would  be  cheaper?     Would  the  cost  of  the 
manufacture  depend  primarily  on  the  cost  of  the  metal  used? 

5.  How  many  pounds  of  ammonium  nitrate  will  be  required  to 
furnish  50  pounds  of  nitrous  oxide?     How  many  liters  of  the  gas 
will  this  give,  counting  453  grams  to  the  pound? 

6.  How  many  pounds  of  aqua  ammonia  of  sp.  gr.  0.90  and  con- 
taining 25  per  cent  of  the  gas  and  how  many  pounds  of  nitric 
acid  of  sp.  gr.  1.42  and  containing  70  per  cent  of  nitric  acid  will  be 
required  to  furnish  the  ammonium  nitrate  for  the  last  problem? 
The  specific  gravities  are  given  as  a  matter  of  information  and  are 
not  to  be  used  in  the  calculation. 


CHAPTER  XIV 


AIR;  THE  NOBLE  GASES;  GROUP  ZERO 

Composition  of  the  Air.  Oxygen. — If  air  is  left  for  some 
time  in  contact  with  phosphorus,  the  oxygen  will  gradually 
combine  with  the  phosphorus  leaving  the  ^^ 

other  gases  contained  in  the  air  unchanged. 

A  quite  accurate  determination  of  the 
per  cent  of  oxygen  in  air  may  be  made  by  ;  ~~| 

the  apparatus  shown  in  Fig.  30.  A  volume 
of  air  is  measured  over  water  in  the  gradu- 
ated tube  (eudiometer),  and  then  a  piece 
of  phosphorus  on  the  end  of  a  wire  is  in- 
serted. After  some  hours  the  phosphorus 
is  removed  and  the  decrease  in  volume 
compared  with  the  original  volume  will  give 
the  per  cent  of  oxygen  in  the  air. 

Composition  of  the  Air.  Carbon  Di- 
oxide.— If  air  is  drawn  through  lime  water 
(a  solution  of  calcium  hydroxide,  Ca(OH)2), 
the  latter  slowly  becomes  turbid  from  the 
formation  of  calcium  carbonate,  CaC03: 

Ca(OH)2  H-  CO2  =  CaCO3  +  H2O 

The  same  test  has  been  applied  to  show 
the  presence  of  carbon  dioxide  in  the  gas 
formed  by  burning  charcoal  in  oxygen  (p. 
10).  The  presence  of  carbon  dioxide  in 
the  breath  can  be  shown  by  blowing  air  from  the  lungs 
through  lime  water,  which  will  quickly  grow  turbid  from 
the  separation  of  calcium  carbonate. 

155 


FIG.  30. 


156  AIR;  THE  NOBLE  GASES;  GROUP  ZERO 

Since  coal,  wood,  oil,  gasolene  and  other  compounds  con- 
taining carbon  are  constantly  burning  all  over  the  world 
and  since  the  breath  of  millions  of  men  and  animals  is  con- 
stantly exhaled  into  the  air,  it  is  evident  that  some  carbon 
dioxide  must  always  be  present  in  the  atmosphere. 

Some  carbon  dioxide  escapes  into  the  air  from  volcanoes. 
Large  quantities  are  formed  from  the  decay  of  organic 
matter  in  the  soil  and  elsewhere  under  the  influence  of 
bacteria. 

If  no  means  were  provided  for  the  removal  of  carbon 
dioxide  from  the  air  it  is  evident  that  in  the  course  of  time 
the  amount  in  the  air  must  very  greatly  increase.  This 
is  prevented  by  the  growth  of  trees  "and  plants,  which  secure 
their  supply  of  carbon  from  the  carbon  dioxide  of  the  air 
and  constantly  return  the  oxygen  of  the  carbon  dioxide  to 
the  air.  The  energy  for  this  process  of  reduction  comes 
from  the  sunlight. 

The  amount  of  carbon  dioxide  normally  present  in  the  air 
is  about  3  parts  in  10,000  or  0.03  per  cent  by  volume. 

Composition  of  the  Air.  Water  Vapor. — The  proportion 
of  oxygen  and  of  carbon  dioxide  in  air  out  of  doors  varies 
very  little  indeed,  but  the  per  cent  of  water  vapor  varies 
between  wide  limits.  In  a  tropical  climate  air  saturated 
with  water  vapor  might  contain  5  per  cent  or  more  by 
volume  of  the  vapor,  while  in  winter  with  the  temperature 
at  10°  below  zero  the  volume  could  not  be  as  much  as  0.3 
per  cent  unless  the  air  were  supersaturated.  Evidence 
for  the  presence  of  water  in  the  air  is  so  common  that  no 
experiments  need  be  given  to  demonstrate  it. 

Composition  of  the  Air.  Argon. — As  long  ago  as  1785 
the  English  chemist,  Cavendish,  mixed  air  with  an  excess 
of  oxygen,  passed  the  sparks  from  a  frictional  electrical 
machine  through  the  mixture  and  absorbed  the  oxides  of 
nitrogen  with  an  alkali.  After  continuing  the  experiment 
for  a  long  time  he  absorbed  the  oxygen  which  was  left  with 


DISCOVERY  OF  ARGON  157 

" liver  of  sulfur"  and  found  that  the  gas  which  remained  was 
not  more  than  Jj[20  °f  ^ne  volume  of  the  original  air.  For 
more  than  100  years  this  experiment  was  taken  as  proof 
that  the  gas  remaining  when  oxygen,  carbon  dioxide  and 
water  vapor  have  been  removed  from  the  air  is  pure  nitrogen. 

In  1894  Lord  Rayleigh,  another  Englishman,  undertook 
to  make  a  very  careful  determination  of  the  weight  of  a 
liter  of  nitrogen.  For  this  purpose  he  prepared  and  weighed 
the  gas  which  remains  when  oxygen,  carbon  dioxide  and 
water  vapor  are  removed  from  common  air.  This  gas  was 
then  supposed  to  be  pure  nitrogen.  For  some  reason  Lord 
Rayleigh  thought  that  it  would  be  of  interest  to  prepare  nitro- 
gen from  ammonia  or  from  some  other  compound  of  nitrogen. 
To  his  surprise  he  found  that  a  liter  of  pure  nitrogen  pre- 
pared by  a  chemical  method  weighs  about  5  milligrams  less 
than  a  liter  of  the  gas  which  he  had  prepared  from  air.  This 
was  a  much  greater  difference  than  could  be  accounted  for 
by  the  ordinary  errors  of  the  determination.  Accordingly 
Lord  Rayleigh,  with  the  assistance  of  Sir  William  Ramsay, 
repeated  the  Cavendish  experiment  and  obtained  a  gas 
which  is  very  much  heavier  than  nitrogen  and  which  they 
called  argon. 

The  Noble  Gases. — Argon  proved  to  be  not  only  a  new 
and  hitherto  unknown  element,  but  an  element  wholly 
different  in  its  properties  from  all  other  elements  then  known. 
The  most  distinctive  property  of  the  element  is  that  it 
seems  to  have  no  chemical  affinity.  It  has  not  been  found 
possible  to  prepare  a  compound  of  argon  with  any  other 
element. 

Within  a  few  years  Sir  William  Ramsay  discovered  five 
other  gaseous  elements  with  similar  properties.  Gold 
and  platinum  have  been  called  noble  metals  because  they 
do  not  tarnish  in  the  air  and  no  single,  ordinary  acid  will 
attack  them.  Because  of  their  lack  of  chemical  affinity 
the  elements  of  this  family  have  been  called,  for  a  similar 


158  AIR;  THE  NOBLE  GASES;  GROUP  ZERO 

reason,  the  noble  gases.     They  are  helium,  neon,  argon, 
krypton,  xenon  and  niton. 

The  relation  of  the  elements  of  this  group  to  the  halogen 
family  on  the  one  side  and  to  the  alkali  metals  on  the  other 
is  shown  in  the  following  table: 

He     4  Li      7 

F   19  Ne    20  Na    23 

Cl  35.5  A     40  K    39 

Br  80  Kr    83  Rb  85.8 

I  127  Xe  130  Cs  133 

Nt  222 

Helium  was  discovered  by  Lockyer  in  the  atmosphere  of 
the  sun  in  1868.  It  was  not  known  that  it  exists  on  the 
earth  till  it  was  discovered  by  Ramsay  in  some  minerals, 
a  few  years  after  the  discovery  of  argon.  Some  years  later 
it  was  shown  that  helium  is  formed  by  the  spontaneous 
decomposition  of  radium  and  of  other  radioactive  elements. 
As  our  knowledge  of  the  composite  nature  of  the  elements 
increases  it  becomes  more  and  more  probable  that  helium 
forms  a  part  of  the  atoms  of  many  different  elements. 

Neon,  Krypton  and  Xenon  are  present  in  very  small 
amounts  in  the  air. 

Niton  is  formed  by  the  decomposition  of  radium,  helium 
being  formed  at  the  same  time.  Niton  in  turn,  decom- 
poses very  much  more  rapidly  than  radium  and  it  is  possi- 
ble to  obtain  only  very  minute  quantities  of  the  gas.  The 
density  was  determined  by  Ramsay  with  the  use  of  only  a 
few  cubic  millimeters.  The  element  was  called  at  first  ra- 
dium emanation.  It  is  powerfully  radioactive  (p.  269). 

Air  is  a  Mixture. — That  the  elements  present  in  air  are  not 
combined  with  each  other  is  evident  from  the  following  facts: 

1.  Wherever  chemical  combination  occurs  heat  is  gen- 
erated or  absorbed.  There  is  no  change  in  temperature 
when  oxygen,  nitrogen  and  argon  are  mixed. 


VENTILATION  159 

2.  If  water  is  boiled,  a  gas  resembling  air  escapes  from  it 
but  when  the  gas  is  examined  it  is  found  that  it  contains  a 
larger  per  cent  of  oxygen  than  air  does.     It  has  been  shown 
that  for  such  gases  as  oxygen  and  nitrogen  the  weight  of  the 
gas  absorbed  by  a  given  volume  of  water  varies  directly 
with  the  pressure  (Henry's  law).     Oxygen  is  more  soluble 
than  nitrogen,  in  water,  and  when  we  take  account  of  the 
solubility  of  the  two  gases  and  of  their  relative  pressures 
in  air  (0,  0.21;  N,  0.78)  the  composition  of  the  gas  expelled 
by  boiling  water  agrees  with  that  which  is  calculated  on  the 
supposition  that  the  oxygen  and  nitrogen  are  simply  mixed 
together.     If  they  were  combined,  the  composition  of  the 
gas  expelled  from  the  water  would  be  the  same  as  the 
composition  of  the  air. 

3 .  The  weight  of  a  liter  of  dry  air  is  almost  exactly  that  cal- 
culated on  the  supposition  that  the  four  gases,  oxygen,  nitro- 
gen, carbon  dioxide  and  argon,  mix  without  combination. 

Weight  of  Proportion  in 

one  liter  the  air 

Oxygen 1.429  X  0.2095     =  0.2994 

Carbon  dioxide ....       1 . 9768  X  0 . 0003     =  0 . 0006 

Argon 1.7828  X  0.0094     =  0.0168 

Nitrogen 1.2507  X  0.7808     =  0.9765 

1.2933 

The  weight  of  a  liter  of  air  by  direct  determination  is 
1.2928  grams. 

Ventilation. — It  was  formerly  supposed  that  the  accumu- 
lation of  carbon  dioxide  from  the  breath  is  the  chief  source 
of  danger  in  poorly  ventilated  rooms.  It  has  now  been 
demonstrated  that  this  is  not  true  and  that  the  amount  of 
carbon  dioxide  present  in  the  air  of  even  badly  ventilated 
rooms  is  never  great  enough  to  cause  any  injury  to  human 
beings. 

In  spite  of  this  it  must  be  considered  as  well  established 
that  lack  of  ventilation  in  factories,  offices  and  dwellings 


160  AIR;  THE  NOBLE  GASES;  GROUP  ZERO 

is  a  frequent  cause  of  disease.  It  is  also  very  well  estab- 
lished that  abundance  of  fresh  air,  secured  by  life  out  of 
doors  both  by  night  and  day,  combined  with  a  nourishing 
diet,  furnishes  the  best  hope  of  recovery  from  incipient 
tuberculosis. 

While  exhaled  carbon  dioxide  is  not  in  itself  harmful,  it 
furnishes  the  best  means  of  determining  whether  a  room  is 
properly  ventilated  or  not.  The  amount  of  the  gas  should 
not  exceed  0.07  per  cent  by  volume.  To  secure  this  stand- 
ard of  ventilation  55,000  liters  or  2000  cubic  feet  of  fresh 
air  will  be  required  each  hour  for  each  person  in  the  room. 

SUMMARY 

The  per  cent  of  oxygen  in  the  air  may  be  determined  by 
absorbing  the  oxygen  with  phosphorus. 

The  air  receives  carbon  dioxide  from  the  breath  of  men 
and  animals,  from  the  combustion  and  decay  of  organic 
matter  and  from  volcanoes. 

Carbon  dioxide  is  removed  from  the  air  by  growing  plants. 
The  amount  in  ordinary  air  is  0.03  per  cent  or  about  3 
parts  in  10,000. 

Water  vapor  in  the  air  varies  between  wide  limits. 

Argon  was  discovered  by  Lord  Rayleigh  in  an  attempt  to 
determine  the  exact  weight  of  a  liter  of  nitrogen. 

The  Zero  Group  of  elements,  sometimes  called  the  noble 
gases,  consists  of  helium,  neon,  argon,  krypton,  xenon  and 
niton. 

Helium  was  discovered  in  the  sun  by  Lockyer.  It  was 
discovered  in  some  minerals  later. 

Niton  is  formed  by  the  disintegration  of  radium. 

Air  is  believed  to  be  a  mixture  because  no  heat  is  evolved 
when  nitrogen  and  oxygen  are  mixed,  because  water  dis- 
solves the  oxygen  and  nitrogen  independently,  and  because 
the  weight  of  a  liter  corresponds  to  the  calculated  weight, 
on  the  supposition  that  it  is  a  mixture. 


EXERCISES.     AIR  161 

Good  ventilation  of  rooms  in  which  people  live,  and  es- 
pecially of  sleeping  rooms,  office  rooms  and  audience  rooms 
is  very  important.  The  carbon  dioxide  in  such  rooms 
should  not  exceed  0.07  per  cent,  though  the  carbon  dioxide 
does  not  seem  to  be,  in  itself,  harmful. 

EXERCISES 

1.  What  is   the   weight   of   a  gram-molecular  volume  of  air? 

2.  How  many  liters  of  air  will  be  required  to  burn  32  grams  of 
sulfur?     31    grams   of  phosphorus?     12   grams  of  carbon?     56 
grams    of   iron? 

3.  What  facts  of  common  experience  demonstrate  the  presence 
of  water  in  the  air? 

4.  What  weight  of  water  vapor  do  22.4  liters  of  air  at  20°  con- 
tain, assuming  that  the  weight  of  a  gram-molecular  volume  of 
water  vapor  is  18  grams  under  normal  conditions?     What  per 
cent  of  water  vapor  does  the  air  contain  by  weight?     Is  moist  air 
heavier  or  lighter  than  dry  air? 

5.  A  sample  of  bituminous  coal  has  the  following  composition: 

Carbon 66 . 25  per  cent 

Hydrogen  (other  than  moisture) ...  4 . 25  per  cent 

Oxygen 8 . 00  per  cent 

Nitrogen 1 . 50  per  cent 

Moisture 9 . 00  per  cent 

Ash 11 .00  per  cent 

How  many  grams  of  oxygen  will  be  required  to  burn  one  pound 
(453  grams)  of  the  coal?  How  many  liters  of  oxygen?  How 
manv  liters  of  air? 


CHAPTER  XV 
THE  PERIODIC  SYSTEM 

When  the  elements  are  arranged  in  the  order  of  their 
atomic  weights  it  is  found  that  elements  of  similar  properties 
occur  at  regular  intervals.  This  fact  has  already  been  used 
as  the  basis  for  the  classification  of  the  non-metallic  elements 
which  have  been  studied  in  the  preceding  pages.  A  classi- 
fication of  all  of  the  elements  on  the  basis  of  their  atomic 
weights  is  given  in  the  accompanying  tables: 

For  an  elementary  study  of  these  tables  the  following 
relations  are  of  most  importance: 

1.  Groups  and  Periods. — The  elements  in  a  perpendicular 
column  are  classed  together  as  a  group.  In  the  first  table 
there  are  nine  groups  numbered  0  to  VIII  but  it.  should 
be  noticed  that  Group  VIII  includes  three  sets  of  elements, 
of  three  in  each  set,  and  that  it  differs  in  this  respect  from 
the  other  groups,  which  contain  only  a  single  element 
at  a  given  point,  in  the  earlier  part  of  the  table. 

The  elements  in  a  horizontal  line  are  classed  together  as 
periods,  but  it  is  necessary,  here,  to  distinguish  between 
the  two  "short  periods"  which  include  the  first  two  lines 
of  the  first  table  and  two  "longer  periods"  which  include 
the  next  four  lines  of  the  first  table,  or  the  third  and  fourth 
lines  of  the  second  table.  The  second  table  is  designed 
to  bring  out  these  long  periods  more  clearly  and  to  show 
that  in  the  first  table  after  chlorine  the  alternate  elements 
of  each  group  resemble  each  other  more  closely  than  the 
successive  elements  in  the  group.  Thus  potassium,  rubid- 
ium and  caesium,  of  Group  I,  are  closely  alike  and  are  usually 

162 


COMPOSITE  NATURE  OF  ATOMS  163 

spoken  of  as  the  first  division  of  the  group,  while  copper, 
silver  and  gold  are  alike  and  form  the  second  division  of 
the  group. 

In  the  rest  of  the  table  the  fifth  and  sixth  periods  are  still 
longer  than  the  third  and  fourth  as  shown  in  the  second 
table.  If  the  rare  earth  elements  are  included,  as  they 
certainly  must  be  if  we  are  not  to  lay  ourselves  open  to  the 
objection  that  facts  have  been  distorted  to  fit  the  theory  of 
the  table,  the  fifth  period  must  include  all  elements  from 
xenon  to  bismuth,  and  only  four  elements,  whose  atomic 
weights  have  been  determined,  are  known  for  the  sixth 
period. 

2.  Composite  Nature  of  Atoms. — Long  before  the  radio- 
active elements  (p.  268)  were  discovered  many  chemists 
considered  that  the  periodic  system  points  clearly  toward 
a  belief  in  the  composite  nature  of  the  atoms  of  the  elements. 
The  discovery  of  the  disintegration  of  the  atoms  of  radio- 
active elements  has  given  almost  conclusive  proof  of  the 
truth  of  this  view  and  has  shown  that  atoms  of  helium  and 
electrons   (atoms  of  negative  electricity)   are  constituent 
parts  of  the  atoms  of  some  of  these  elements.     For  a  long 
time  some  chemists  have  proposed  the  hypothesis  that 
atoms  of  hydrogen  are  also  constituent  parts  of  the  atoms 
of  the  other  elements,  and  some  facts  point  rather  strongly 
toward  such  a  view.     If  we  suppose,  as  seems  reasonable, 
that  all  atoms  are  composite,  in  order  to  account  for  the 
atomic  weights  it  seems  necessary  to  suppose  that  particles 
which  are  intermediate  in  weight  between  helium  (4)  and 
electrons   (1/1800  the  weight  of  hydrogen  atoms)   must 
enter  into  the  composition. 

Thus  far,  however,  no  one  has  shown  that  hydrogen 
atoms  are  found  among  the  disintegration  products  of  any 
element. 

3.  The  position  of  an  element  in  the  system  indicates  at 
least  one  of  its  valences,  especially  its  highest  valence  toward 


164 


THE  PERIODIC  SYSTEM 


0 

C 


s? 


fctf 

lo. 
oS 


II 


10 


PQ  ^H 


o 

§tf 

u 


- 


COMPOSITE  NATURE  OF  ATOMS 


165 


i,  SS 


ff 


S     0 


58     OT 

o  g 


ss 


s 


Ml 


c,  a 


a  * 


H 


rt   o 
0   o 


166  THE  PERIODIC  SYSTEM 

oxygen  (positive  valence)  and,  for  the  non-metallic  ele- 
ments, the  valence  toward  hydrogen  (negative  valence). 
This  will  be  clear  from  the  following  formulas  of  compounds 
of  the  elements  of  the  second  period. 

Group  I        Group  II      Group  III      Group  IV      Group  V      Group  VI       Group  VII 

Na  Mg=O  O  O  O  O         O    O 

\  /"          ./        /"          /         II  / 

Cl=( 


O                        Al  C            N=O  S=O          C1=O 

/                               \  \            \  \                 \ 

Na                                    O  O            O  O                 O 

Al  N=O  C1=O 

\  \  II  \ 

O  O  O     O 


H    H  H  H 

N— H        S  Cl— H 


i-H 


H  H  H 

It  should  be  noticed  that  the  sum  of  the  valences  toward 
oxygen  and  toward  hydrogen  is  eight.  This  is  doubtless 
connected  in  some  way  with  the  structure  of  the  atoms,  but 
our  knowledge  of  that  structure  is  still  too  indefinite  to  give 
any  indication  of  the  basis  for  this  or,  indeed,  for  many 
other  puzzling  facts. 

4.  In  the  periods  metallic  properties  decrease  and  non- 
metallic  properties  increase  from  left  to  right,  that  is, 
with  increasing  atomic  weight.  In  the  groups,  on  the 
other  hand,  metallic  properties  increase  and  non-metallic 
properties  decrease  from  top  to  bottom.  The  result  is 
that  the  non-metallic  elements  are  found  in  the  upper  right- 
hand  corner  of  the  table,  above  the  dotted  line.  Strictly 
speaking,  the  Zero  Group  has  no  chemical  properties  and 
we  cannot  classify  the  elements  of  that  group  as  either 
metallic  or  non-metallic.  It  would,  perhaps,  be  more 
logical  to  place  the  group  to  the  right  of  the  table  as  alter- 
nate with  the  elements  of  Group  VIII.  With  such  an 


g      3 


168  THE  PERIODIC  SYSTEM 

arrangement  there  would  be  two  divisions  in  Group  VIII, 
as  in  the  other  groups. 

5.  The   physical   properties   of   the   elements,    such   as 
melting  points  and  specific  gravity,  are  closely  connected 
with  their  positions  in  the  table.     The  relation  between 
the  melting  points  and  the  positions  in  the  table  will  be 
seen  by  an  examination  of  the  melting  points  given  in  the 
second  table.     The  use  of  tungsten  (W)  for  electric  lights 
was  suggested  by  the  position  of  the  element  in  the  tab'le. 

The  relation  of  the  specific  gravity  of  the  elements  to 
their  positions  in  the  table  is  shown  in  Fig.  31.  The 
atomic  volume  of  an  element  is  the  volume  occupied  by  one 
gram  atom.  Thus  the  specific  gravity  of  potassium  (K) 
is  0.862  and  a  gram  atom  of  the  element  weighs  39.1  grams. 

39  1 
The  volume  occupied  by  a  gram  atom  will  be  Q  g62  = 

45.4  cc.  and  this  is  the  atomic  volume  of  potassium. 

6.  Atomic  Numbers. — An  examination  of  the   periodic 
table  reveals  the   fact   that   three   elements,   argon    (A), 
cobalt  (Co)  and  tellurium  (Te),  are  not  placed  in  the  table 
in  accordance  with  their  atomic  weights.     As  the  properties 
of  these  elements  leave  no  question  but  that  they  are  prop- 
erly placed,  these  exceptions  have  caused  some  misgiving 
as  to  the  validity  of  the  table.     This  has  led  to  a  number  of 
very  careful  determinations  of  the  atomic  weights  of  these 
elements  and  especially  of  tellurium.     These  determinations 
have  confirmed  the  values  for  the  atomic  weights  which  are 
given  and  facts  of  this  sort  are  always  accepted  by  chemists 
no  matter  how  much  they  may  conflict  with  theories  or 
systems  which  are  in  vogue. 

In  comparatively  recent  times  it  has  been  discovered  by 
means  of  X-ray  spectra  that  each  element  has  a  charac- 
teristic property,  called  its  atomic  number.  This  atomic 
number  is  approximately  one-half  of  the  atomic  weight  and 
appears  to  be  connected  with  the  structure  of  the  atom  in 


EXERCISES  169 

such  a  manner  that  it  is  a  more  fundamental  charac- 
istic  than  the  atomic  weight.  The  atomic  numbers  of 
argon,  cobalt  and  tellurium  agree  with  the  positions  assigned 
them  in  the  periodic  table. 

SUMMARY 

Elements  are  classified  in  a  table,  called  the  periodic 
system,  which  is  based  on  bheir  atomic  weights. 

There  are  nine  groups  of  elements  and  each  group  except 
the  Zero  Group  and  Group  VIII  is  separated  into  two  divi- 
sions. 

There  are  two  " short  periods,"  two  "long  periods," 
a  fifth  very  long  period  and  a  sixth  period  for  which  only  a 
very  few  elements  are  known. 

The  valences,  metallic  or  non-metallic  properties,  melting 
points  and  specific  gravities  of  the  elements  are1  closely 
connected  with  the  positions  of  the  elements  in  the  periodic 
system. 

The  atomic  volume  of  an  element  is  the  volume  in  cubic 
centimeters  occupied  by  one  gram  atom  of  the  element. 

The  atomic  numbers  of  the  elements,"  determined  by 
means  of  their  X-ray  spectra,  are  even  more  characteristic 
than  their  atomic  weights. 

EXERCISES 

1.  Prepare  tables  of  the  oxides,  acids,  chlorides  and  hydrogen 
compounds  of  the  elements  which  have  been  studied,  showing 
the  relationships  between  them. 

2.  The  specific  gravity  of  copper  is  8.93.     What  is  its  atomic 
volume? 

3.  From  the  positions  in  the  periodic  system  what  oxides  are  to 
be  expected  from  the  following  elements:    Strontium,  Sr,  tita- 
nium, Ti,  vanadium,  V,  chromium,  Cr,  molybdenum,  Mo,  thor- 
ium, Th,  tantalum,  Ta? 

4.  What  elements  besides  carbon  and  tungsten  have  been  used 
for  filaments  of  electric  lights? 


CHAPTER  XVI 

GROUP  V:  PHOSPHORUS,  ARSENIC,  ANTIMONY  AND 

BISMUTH 

The  Nitrogen  Family  of  Elements. — Fluorine  differs 
very  decidedly  in  many  of  its  properties  from  the  other- 
members  of  the  halogen  family,  chlorine,  bromine  and  iodine, 
and  oxygen  is  quite  different  from  sulfur,  selenium  and 
tellurium.  In  a  similar  manner  nitrogen  differs  from  phos- 
phorus, arsenic,  antimony  and  bismuth.  In  spite  of  these 
differences,  however,  there  are  enough  resemblances  so 
that  a  comparison  of  the  compounds  and  properties  is  of 
much  service  in  studying  the  elements  of  the  group. 

Occurrence  of  Phosphorus. — Phosphorus  is  not  found 
free  in  nature.  While  free  nitrogen  is  induced  to  combine 
with  other  elements  with  difficulty  and  only  under  very 
special  conditions,  ordinary  phosphorus  combines  with 
oxygen  so  readily  and  holds  the  element  so  firmly  that  it  is 
impossible  that  it  should  be  liberated  or  exist  free  under  any 
ordinary  conditions  found  in  nature.  It  is  found  almost  or 
quite  exclusively  in  the  form  of  salts  of  phosphoric  acid, 
H3PO4.  The  most  common  salt  is  tricalcium  phosphate, 
Ca3(PO4)2.  This  salt  is  the  principal  mineral  constituent 
of  bones  and  so  of  bone-ash,  obtained  by  burning  bones.  It 
is  also  found  as  a  mineral,  more  or  less  pure,  in  North  and 
South  Carolina,  Tennessee,  Georgia,  Florida  and  in  some 
other  states.  Since  phosphates  are  essential  for  the  growth 
of  plants  and  some  soils  are  deficient  in  them,  large  quan- 
tities of  these  phosphates  are  used  for  fertilizing  purposes. 
Apatite  is  another  mineral  containing  phosphorus.  It  is 
a  combination  of  calcium  phosphate  with  the  fluoride, 

170 


PREPARATION  AND  PROPERTIES  OF  PHOSPHORUS     171 

CaF2,  or  chloride,  CaCl2,  and  has  the  formula  Ca5  (PO4)3F 
or  Ca5(PO4)3Cl.  The  relation  between  the  valence  of 
calcium  and  those  of  the  phosphate  radical,  P04,  and  of 
fluorine  should  be  noticed  in  this  formula. 

Preparation  of  Phosphorus. — Phosphorus  has  been  used 
in  the  manufacture  of  matches  for  less  than  a  century 
but  during  that  time  the  manufacture  of  the  element  has  be- 
come very  important.  For  a  long  time  it  was  obtained  by 
a  complicated,  tedious  process  in  which  calcium  phosphate 
was  first  treated  with  sulfuric  acid  and  the  phosphoric 
acid  or  acid  calcium  phosphate  formed  was  reduced  by 
heating  it  with  carbon.  During  comparatively  recent  times 
the  high  temperature  which  can  be  secured  in  an  electric 
furnace  has  made  it  possible  to  manufacture  phosphorus  by 
heating  a  mixture  of  calcium  phosphate,  Ca3(PO4)2,  silica, 
or  sand,  SiO2,  and  carbon,  in  the  form  of  charcoal  or  coke: 

2Ca3(PO4)2  +  6SiO2  +  IOC  =  6CaSiO3  +  P4  +  10CO 

The  silica,  SiO2,  combines  with  the  calcium  and  oxygen 
(lime),  CaO,  to  form  calcium  silicate,  CaSiO3,  while  the 
carbon  takes  oxygen  from  the  phosphorus.  In  writing  the 
equation  notice  that  the  formula  of  phosphorus  is  P4  and 
that  two  molecules  of  calcium  phosphate  are  required  to 
give  this.  The  rest  follows  logically. 

Properties  of  Phosphorus. — Ordinary  phosphorus  is  a 
wax-like,  almost  colorless  solid  when  first  prepared,  but  it 
soon  acquires  a  yellow  or  red  color  on  exposure  to  the  light. 
Its  kindling  temperature  is  very  low  and  it  is  kept  under 
water  and  cast  into  sticks.  It  boils  at  290°  and  may  be 
easily  distilled.  Ordinary  phosphorus  is  very  poisonous. 

Red  Phosphorus. — If  heated  at  240°-250°  for  some  time 
in  a  closed '  vessel  ordinary  phosphorus  is  changed  partly, 
but  not  completely,  to  the  allotropic  form  known  as  red 
phosphorus.  At  a  higher  temperature  red  phosphorus  is 
converted  into  a  vapor  which  condenses,  on  cooling,  to  the 


172  PHOSPHORUS,  ARSENIC,  ETC. 

ordinary  form.  Red  phosphorus  does  not  take  fire  with 
ordinary  friction  and  it  is  not  poisonous.  It  is  used  in  the 
manufacture  of  safety  matches  and  for  many  purposes  in 
chemical  laboratories. 

Matches. — Phosphorus  was  discovered  nearly  250  years 
ago  but  it  was  not  used  for  the  manufacture  of  matches  till 
1827.  Persons  still  living  can  remember  being  sent  to  a 
neighbor's  to  borrow  coals  of  fire  because  of  the  difficulty 
of  starting  a  fire  with  flint,  steel  and  tinder.  During  the 
nineteenth  century  ordinary  phosphorus  was  chiefly  used 
for  making  matches,  though  in  European  countries  the  or- 
dinary phosphorus  match  came  to  be  displaced  by  safety 
matches,  which  must  be  ignited  on  a  prepared  surface  and 
offer  less  danger  of  accidental  fires.  The  use  of  ordinary 
matches  in  America  is  one  of  the  reasons  for  the  very  much 
greater  losses  by  fire  here  than  abroad. 

Owing  to  the  mortality  due  to  phosphorus  poisoning  in 
match  factories,  Congress  in  1912  passed  a  law  taxing 
the  use  of  ordinary  phosphorus,  and  tetraphosphorus 
trisulfide,  P^Ss,  is  now  used  in  place  of  it.  This  gives  a 
match  which  can  be  ignited  by  friction  and  still  its  manu- 
facture does  not  endanger  the  health  and  lives  of  the  work- 
men. The  match  sticks  are  made  of  some  wood  which 
does  not  readily  retain  a  live  coal  after  the  flame  is  ex- 
tinguished and  this  non-inflammable  quality  is  further 
increased  by  treatment  with  a  dilute  solution  of  magnesium 
sulfate  or  alum. 

Phosphine. — When  phosphorus  is  warmed  with  a  con- 
centrated solution  of  sodium  hydroxide  (Fig.  32),  NaOH, 
phosphine,  PH3,  a  compound  corresponding  to  ammonia, 
NH3,  is  formed.  The  details  of  the  reaction  need  not  be 
given  here.  A  small  amount  of  a  second  compound  of 
phosphorus  and  hydrogen,  P2H4,  is  also  formed  and  the 
presence  of  this  causes  the  phosphine  prepared  by  the 
method  described  to  take  fire  on  coming  to  the  air. 


PHOSPHORUS  PENTOXIDE 


173 


Phosphine  combines  with  hydriodic  acid,  HI,  to  form  the 
salt  phosphonium  iodide,  PH4I,  which  corresponds  to  the 
similar  salts  ammonium  iodide,  NH4I,  and  ammonium 
chloride,  NH4C1,  formed  from  ammonia.  Phosphonium 
iodide  is  far  less  stable  than  ammonium  iodide,  and  the 
compound  is  chiefly  of  interest  because  of  the  analogy 
which  it  shows  between  phosphine  and  ammonia. 


FIG.  32. 

Phosphorus  Pentoxide,  or  Phosphoric  Anhydride. — When 
phosphorus  is  burned  in  dry  air  or  in  oxygen  the  oxide 
formed  has  the  composition  represented  by  the  formula 
P2O5.  The  oxide  may  be  converted  into  a  vapor  at  a  rather 
high  temperature  and  it  has  been  found  that  the  gram- 
molecular  volume  of  the  vapor  weighs  about  284  grams, 
from  which  it  follows  that  the  true  formula  is  P4Oio  and 
not  P2O5.  The  name  phosphorus  pentoxide  is  still  used, 
however,  and  the  formula  P2O5  correctly  represents  the 
composition  of  the  compound,  of  course.  The  name  phos- 
phoric anhydride  is  strictly  accurate  and  is,  perhaps,  better. 
It  is  usually  more  convenient  to  use  the  formula  P2Os  in 


174  PHOSPHORUS,  ARSENIC,  ETC. 

equations,  just  as  equations  are  sometimes  written  in  which 
oxygen,  hydrogen  and  other  elements  are  involved  without 
taking  account  of  the  true  formulas,  02,  H2,  etc. 

Phosphorus  pentoxide  has  a  very  strong  affinity  for  water 
and  it  is  the  most  perfect  drying  agent  we  have  for  the 
gases  which  do  not  combine  with  it. 

Phosphoric  Acids. — Phosphorus  pentoxide  may  combine 
with  one,  two  or  three  molecules  of  water  to  form  three  dif- 
ferent acids,  which  are  all  called  phosphoric  acids  because 
the  phosphorus  is  in  the  same  state  of  oxidation  in  each. 
The  addition  or  removal  of  water  from  a  compound  is 
considered  neither  as  an  oxidation  nor  a  reduction.  The 
relation  of  the  three  acids  to  phosphorus  pentoxide  is  as 
follows : 

H2OP2O5  =  2HPO3,  Metaphosphoric  acid 
2H2OP2O5  =  H4P207,  Pyrophosphoric  acid 
3H2OP205  =  2H3P04,  Orthophosphoric  acid 

The  last  of  these  acids,  Orthophosphoric  acid,  is  much 
more  common  than  either  of  the  others  and  is  always  under- 
stood when  the  name  phosphoric  acid,  without  a  prefix, 
is  used. 

Salts  of  Orthophosphoric  Acid. — As  a  tribasic  acid 
Orthophosphoric  acid  forms  three  classes  of  salts:  acid 
salts  (p.  76),  in  which  one  or  two  of  the  hydrogen  atoms 
of  the  acid  have  been  replaced,  and  normal  salts,  in  which 
all  three  atoms  have  been  replaced  by  a  metal.  The  salts 
are  distinguished  as  primary,  secondary  and  tertiary,  or, 
more  often,  by  the  prefixes  mono-,  di-  and  tri-.  The 
latter  refer  to  the  numbers  of  hydrogen  atoms  replaced 
from  one  molecule  of  the  acid.  Thus  Ca(H2PO4)2  or  CaH4- 
(PO4)2  is  monocalcium  phosphate.  CaHP04  is  dicalcium 
phosphate,  although  it  has  only  a  single  atom  of  calcium 
in  the  molecule. 


HYDROLYSIS  OF  SALTS  OF  PHOSPHORIC  ACID     175 
The  relations  will  be  clear  from  the  following  table: 


Mpnosodium  phosphate,  NaH2PO4        Monocalcium  phosphate, 
Disodium  phosphate,  Na2HPO4  Dicalcium  phosphate,  CaHPO4 

Trisodium  phosphate,  NasPO4  Tricalcium  phosphate,  Cas(PO4)z 

Salts  of  pyrophosphoric  acid  and  of  metaphosphoric  acid 
are  formed  by  heating  the  acid  salts  or  mixed  ammonium 
salts  derived  from  orthophosphoric  acid.  Thus  either 
monosodium  phosphate,  NaH2PO4,  or  sodium  ammonium 
phosphate,  NaNH4HP04,  gives  sodium  metaphosphate, 
NaPO3,  when  heated.  The  sodium  ammonium  phosphate 
loses  both  ammonia  and  water.  Give  the  formulas  and 
names  of  the  salts  which  will  be  formed  by  heating  the 
following:  CaHP04,  MgNH4PO4,  KNH4HPO4,CaH4(P04)2 

Hydrolysis  of  Salts  of  Phosphoric  Acid.  —  Phosphoric 
acid,  H3PO4,  is  a  comparatively  strong  acid.  In  other 
words,  it  is  largely  ionized  in  dilute  solutions,  giving  hydro- 
gen ions,  H+,  and  monophosphate  ions,  H2PO4~.  To  a 
very  slight  extent  it  is  further  ionized  to  two  hydrogen  ions, 
H+,  H+,  and  the  diphosphate  ions,  HPO4=.  The  diphos- 
phate  ions,  HPO4=,  also  ionize  to  a  very  trifling  extent  into 
hydrogen  ions  and  triphosphate  ions,  PO4~,  but  this  last 
ionization  is  probably  less  in  amount  than  even  the  ioniza- 
tion  of  water  to  hydrogen,  H+,  and  hydroxide,  OH~,  ions. 

If  a  solution  of  sodium  hydroxide,  NaOH,  is  added  to  a 
dilute  solution  of  phosphoric  acid,  H3PO4,  in  the  proportion 
of  one  molecule  of  the  hydroxide  for  one  molecule  of  the 
acid,  the  hydrogen  ions  for  the  first  ionization  of  the  acid 
will  be  removed  by  combination  with  the  hydroxide  ions  of 
the  base.  The  solution  will  then  contain  sodium  ions, 
Na+,  'monophosphate  ions,  H2PO4~,  and  a  few  hydrogen  ions 
H+,  and  diphosphate  ions,  HPO4=.  The  solution  will  still 
be  faintly  acid,  but  if  more  sodium  hydroxide  is  added  a 
point  is  soon  reached  where  the  tendency  of  the  diphos- 
phate ions,  HPO4=,  to  combine  with  hydrogen  ions  to  form 
monophosphate  ions,  H2PO4~,  will  just  balance  the  tend- 


176  PHOSPHORUS,  ARSENIC,  ETC. 

ency  of  water  to  separate  into  hydrogen  and  hydroxide 
ions  and  the  numbers  of  hydrogen  and  of  hydroxide  ions 
in  the  solution  will  be  equal;  the  solution  will  then  be 
neutral.  This  point  is  reached  long  before  the  second 
hydrogen  atom  of  the  phosphoric  acid  has  been  neutral- 
ized. Any  further  addition  of  sodium  hydroxide  must 
give  an  excess  of  hydroxide  ions  and  the  solution  will  have 
an  alkaline  reaction. 

If  the  compound  disodium  phosphate,  Na2HPO4,  is 
dissolved  in  water  it  separates  into  sodium  ions,  Na+,  Na+, 
and  diphosphate  ions,  HPO4=.  The  diphosphate  ions  have 
so  strong  a  tendency  to  combine  with  hydrogen  ions,  how- 
ever, that  when  there  are  many  of  them  in  a  solution  they 
will  combine  with  the  hydrogen  ions  of  the  water,  leaving 
an  excess  of  hydroxide  ions.  This  will  give  the  solution  an 
alkaline  reaction-. 

A  decomposition  of  a  salt  or  compound  by  the  ions  of 
water,  in  this  way,  is  called  hydrolysis  (from  i55wp,  water, 
and  Xuco,  I  loose)  because  it  is  caused  by  the  loosening 
effect  of  the  water. 

Other  Acids  of  Phosphorus. — Two  other  acids  of  phos- 
phorus, phosphorous  acid,  H3PO3,  and  hypophosphorous 
acid,  H3PO2,  are  fairly  well  known.  Phosphorous  acid 
is  bibasic.  No  salts  have  been  prepared  in  which  all  three 
of  its  hydrogen  atoms  have  been  replaced.  Hypophos- 
phorous acid  is  monobasic.  What  will  be  the  formulas  of 
the  normal  sodium  and  calcium  salts  of  these  acids? 

Chlorides  of  Phosphorus.  Hydrolysis. — When  phos- 
phorus burns  in  chlorine  either  phosphorus  trichloride, 
PG13,  or  phosphorus  pentachloride,  PC15,  may  be  formed 
according  to  the  conditions  of  the  experiment.  These 
chlorides  are  decomposed  (hydrolyzed)  by  water  with  the 
formation  of  hydrochloric  acid  and  either  phosphorous 
acid  or  phosphoric  acid: 


ARSENIC  177 

Cl      H  -  OH         HO 
P— Cl  +  H  -  OH   =    HO— P  +  3HC1 

Cl      H  -  OH         HO7 

Cl      H-O-H      H-Ov 
/Cl      H-0-H=H-  0-^p  =  o  +  5HC1 
P/.C1      H-O-H      H-O/ 
\SC1      H-O 

H/ 


or  PC13  -f  3H2O  =  H3PO3  +  3HC1 

and  PC15  +  4H2O  =  H3PO4  +  5HC1 

Arsenic,  As,  74.96. — 'Arsenic  is  present  in  considerable 
amounts  in  some  of  the  copper  ores  in  the  West  and  in 
smelting  these  ores  the  arsenic  is  oxidized  to  arsenic  trioxide, 
which  is  produced  in  enormous  quantities  at  the  same 
time.  The  deposit  of  the  arsenic  trioxide  on  vegetation  in 
the  neighborhood  of  the  smelters  has  caused  considerable 
trouble  from  the  poisoning  of  cattle.  Arsenic  is  also  found 
in  arsenical  pyrites,  a  mineral  called  mispickel  and  having  the 
formula  FeAsS.  Arsenic  may  be  obtained  either  by  heating 
mispickel  or  by  reducing  the  trioxide  with  carbon.  When 
not  tarnished  arsenic  is  a  steel-gray,  brittle  metal  which 
may  be  easily  powdered.  When  heated  it  volatilizes  with- 
out melting.  In  the  air  the  vapors  burn  to  the  trioxide, 
As2O3. 

Arsenic  trioxide  is  a  white,  crystalline  powder,  formed  by 
burning  the  metal.  It  is  nearly  or  quite  tasteless  and  is  a 
violent  poison.  It  is  used  as  a  ratsbane  and  has  often  been 
used  in  cases  of  criminal  poisoning.  It  is  also  used  as  a 
medicine.  The  methods  of  chemical  analysis  make  it 
possible  to  detect  a  very  minute  quantity  of  the  element. 

Arsine. — If  a  soluble  compound  of  arsenic  is  introduced 
into  a  generator  in  which  hydrogen  is  prepared  from  zinc 
and  hydrochloric  or  sulfuric  acid,  the  arsenic  is  reduced  to 
arsine,  AsH3.  The  arsine  escapes  with  the  hydrogen.  It 

12 


178 


PHOSPHORUS,  ARSENIC,  ETC. 


imparts  to  the  hydrogen  flame  a  pale  blue  color  and  metallic 
arsenic  is  deposited  on  a  piece  of  cold  porcelain  held  in 
the  flame,  very  much  as  carbon  is  deposited  from  a  candle 
flame.  If  the  hydrogen  containing  arsine  is  passed  through 
a  heated  tube,  the  arsine  is  decomposed  and  a  mirror  of 
arsenic  is  deposited  on  the  walls  of  the  tube.  By  the  arrange- 
ment shown  in  Fig.  33  it  is  possible  to  detect  1/1000  of  a 
milligram  of  arsenic.  It  is,  of  course,  necessary  to  use  zinc 
and  acid,  which  are  entirely  free  from  arsenic. 


FIG.  33. 

Arsenic  trichloride  is  a  colorless  liquid  which  may  be  pre- 
pared by  the  direct  union  of  arsenic  and  chlorine.  It  is 
hydrolyzed  by  water  in  the  same  way  as  phosphorus  tri- 
chloride and  other  chlorides  of  non-metallic  elements. 
Arsenic  trioxide  separates  from  the  solution.  Some  arse- 
nic remains  in  the  solution,  however,  and  the  conduct  of 
the  solution  toward  hydrogen  sulfide  indicates  that  a  part 
of  this  is  still  in  the  form  of  the  chloride.  In  the  reactions 
represented  by  the  equations: 

2AsCl8  +  6HOH  <=*  6HC1  +  2As(OH)3  <=±  As2O3  +  3H2O 

the  equilibrium  is  far  to  the  right  and  the  course  of  the 
reaction  in  that  direction  is  also  favored  by  the  slight  solu- 
bility of  the  arsenic  trioxide. 

Arsenious  Acid. — Just  as  nitrous  acid,  HNO2,  easily 
decomposes  into  nitrous  anhydride,  N2Os,  and  water,  arse- 
nious  acid,  HsAsOa,  decomposes  into  arsenic  trioxide, 


ARSENIC  ACID  179 

As203,  the  anhydride  of  the  acid,  and  water.  Salts  of  the 
acid  are  known,  however,  such  as  silver  arsenite,  Ag3AsO3, 
and  Paris  green,  Cu(C2H3O2)2 .  Cu3(As03)2.  The  latter  is 
a  double  salt  of  copper  acetate  and  copper  arsenite.  It  is 
used  as  a  green  pigment  for  painting  and  as  a  poison  for 
potato  bugs  and  other  insects. 

Arsenic  Acid. — When  arsenious  oxide  is  warmed  with 
nitric  acid  it  is  oxidized  to  arsenic  acid,  H3As04,  while  the 
nitric  acid  is  reduced  to  nitrous  anhydride,  N2O3.  This  is, 
indeed,  the  most  -convenient  method  of  preparing  nitrous 
anhydride : 

As2O3  +  2HNO3  +  2H2O  =  2H3AsO4  +  N2O3 

Arsenic  acid  is  tribasic.  The  two  most  interesting  salts 
are  trisilver  arsemate,  Ag3AsO4,  and  magnesium  ammonium 
arsenate,  MgNH4As04.  What  will  be  the  formula  and 
what  is  the  name  of  the  compound  formed  by  heating  the 
latter  (see  p.  175)? 

Arsenic  trisulfide,  As2S3,  is  formed  by  passing  hydrogen 
sulfide  into  a  solution  of  arsenic  trioxide,  As2O3,  in  hydro- 
chloric acid.  The  formation  of  the  compound  indicates 
the  presence  of  arsenic  trichloride,  AsCl3,  in  the  solution. 
Arsenic  trisulfide  is  found  as  a  mineral  called  orpiment,  and 
is  sometimes  used  as  a  yellow  pigment  by  artists. 

Arsenic  trisulfide  dissolves  easily  in  a  solution  of  am- 
monium sulfide  (NH4)2S,  giving  ammonium  sulf arsenite, 
(NH4)3AsS3: 

As2S3  +  3(NH4)2S  =  2(NH4)3AsS3 

Ammonium  sulf  arsenite  receives  its  name  from  the  simi- 
larity to  ammonium  arsenite  (NH4)3AsO3.  It  is  used  in 
qualitative  analysis  to  separate  arsenic  from  copper,  lead  and 
other  metals  whose  sulfides  do  not  form  similar  compounds. 

Antimony,  Sb,  120.2. — Antimony  occurs  chiefly  as  the 
mineral  stibnite,  which  is  antimony  trisulfide,  Sb2S3. 


180  PHOSPHORUS,  ARSENIC,  ETC. 

Precipitated  antimony  trisulfide  is  orange  colored  but  the 
mineral  stibnite  is  black.  Metallic  antimony  is  obtained 
from  stibnite  by  roasting  it  in  the  air  to  convert  the  metal 
into  the  oxide  and  then  reducing  the  oxide  with  carbon: 

2Sb2S3  +  9O2  =  2Sb2O3  +  6SO2 
2Sb2O3  +  3C    =  4Sb  +  3C02 

Antimony  is  a  brittle  metal  which  may  be  easily  pulver- 
ized. When  heated  in  the  flame  of  a  blowpipe  on  charcoal 
it  melts  easily  and  burns  slowly  to  antimony  trioxide, 
Sb2O3,  which  escapes  as  a  white  cloud. 

In  its  metallic  luster  and  conductivity  for  electricity 
antimony  is  clearly  metallic.  In  its  brittleness  it  resembles 
the  non-metals.  It  is  sometimes  called  a  half-metal. 

An  alloy  containing  lead  and  antimony  is  used  for  type 
metal.  The  antimony  gives  hardness  and  also  causes  the 
metal  to  expand  in  solidifying,  giving  sharp,  clear-cut  type. 
An  alloy  with  lead,  copper  and  a  little  bismuth  is  made  as  an 
antifriction  metal  for  bearings  in  machinery.  One  of  these 
alloys  is  called  Babbitt  metal. 

Stibine,  SbH3,  may  be  prepared  in  the  same  manner  as 
arsine,  AsH3,  and  resembles  it  closely  in  its  properties.  It 
decomposes  at  a  lower  temperature  than  arsine  and  gives 
a  blacker,  more  sooty  spot  on  porcelain.  It  is  to  be  noticed 
that  neither  arsine  or  stibine  combines  with  acids,  differing 
in  this  respect  from  ammonia  and  phosphine. 

Antimony  trichloride,  SbCl3,  may  be  prepared  by  dis- 
solving either  the  trioxide,  Sb2O3,  or  the  trisulfide,  Sb2S3, 
in  concentrated  hydrochloric  acid.  The  solution  may  be 
evaporated  and  the  trichloride  distilled.  This  is  in  very 
marked  contrast  with  the  conduct  of  arsenic  trichloride  and 
illustrates  the  increase  in  metallic  and  decrease  in  non- 
metallic  properties  with  increasing  atomic  weight  in  the 
group.  Chlorides  of  non-metals  are  hydrolyzed  by  water. 
Chlorides  of  metals  dissolve  and  ionize  in  water. 


BISMUTH  181 

Antimony  trichloride,  however,  shows  also  the  properties 
of  a  non-metallic  chloride.  If  the  solution  is  diluted  it  is 
hydrolyzed  in  part  and  antimony  oxychloride,  SbOCl,  is 
precipitated : 

SbCl3  +  H2O  =  SbOCl  +  2HC1 

Oxides  of  Antimony. — Antimony  forms  three  oxides, 
Sb2O3,  Sb2O4  and  Sb2O5,  which  correspond  to  the  three 
highest  oxides  of  nitrogen. 

Antimony  Hydroxide,  Sb(OH)3  or  Antimonious  acid, 
H3SbO3,  may  act  either  as  a  base  or  an  acid: 

Sb(OH)3  +  3HC1  =  SbCl3  +  3H20 
H3SbO3  +  NaOH  =  NaH2SbO3  +  H2O 

Sodium  antimonite 

Here,  again,  it  is  seen  that  the  conduct  of  the  compounds 
of  antimony  places  it  midway  between  the  metals  and  the 
non-metals  and  that  it  is  partly  both  a  metal  and  a  non- 
metal. 

Antimony  trisulfide  is  formed  as  an  orange-colored  pre- 
cipitate when  hydrogen  sulfide  is  passed  into  an  acid  solution 
of  an  antimony  salt.  It  dissolves  in  ammonium  sulfide, 
forming  ammonium  sulfantimonite,  (NH4)3SbS3,  corre- 
sponding to  arsenic  sulfarsenite,  (NH4)3AsS3. 

Bismuth,  Bi,  208. — Bismuth  is  found  mostly  in  the  free 
state  in  nature.  It  is  a  brittle  metal  and  has  a  slightly 
reddish  cast.  It  is  used  in  a  variety  of  alloys,  chiefly  to 
lower  the  melting  point.  Thus  it  is  used  in  the  stereotype 
metal,  employed  in  printing  daily  papers,  to  lower  the 
melting  point  so  far  that  it  does  not  injure  the  papier- 
mache  mold.  Wood's  metal  (Bi  4  parts,  Pb  2  parts,  Sn 
1  part,  Cd  1  part)  melts  at  60.5°.  A  teaspoon  made  from 
the  metal  will  melt  in  a  hot  cup  of  tea.  Alloys  similar 
to  Wood's  metal  are  used  for  safety  plugs  in  steam  boilers 
and  in  automatic  sprinklers  for  protection  against  fire  and 
for  the  safety-fuses  in  electrical  work.  A  small  amount 


182  PHOSPHORUS,  ARSENIC,  ETC. 

of  bismuth  has  been  found  to  improve  the  quality  of  Babbitt 
metal. 

Compounds  of  Bismuth. — Bismuth  forms  no  compound 
corresponding  to  arsine,  AsH3,  and  stibine,  SbH3.  Bismuth 
trioxide,  Bi2O3,  does  not  dissolve  in  alkalies,  and  bismuth 
sulfide,  Bi2S3,  does  not  dissolve  in  ammonium  sulfide. 
In  these  regards  bismuth  is  more  distinctly  metallic  than 
any  other  member  of  the  group.  The  chloride,  BiCl3, 
is  hydrolyzed  by  water  to  the  oxychloride,  BiOCl,  and 
the  nitrate,  Bi(NO)3,  to  the  basic  nitrate,  Bi(OH)2N03,  or 
to  bismuthyl  nitrate,  BiON03: 

/OH 

Bi(NO3)3  +  2HOH  =  Bi— OH  +  2HNO3 


/OH  Q 

Bi— OH  =  Bi/        +  H20 
\N03          XN°3 

This  basic  nitrate,  which  varies  in  composition  according 
to  the  method  of  preparation,  is  used  in  medicine  and  as  a 
face  powder.  Pharmacists  still  often  use  for  it  the  anti- 
quated name  "subnitrate  of  bismuth." 

SUMMARY 

Nitrogen  is  quite  different  from  the  other  elements  of 
Group  V,  phosphorus,  arsenic,  antimony  and  bismuth. 

Phosphorus  is  found  as  tricalcium  phosphate  in  bone- 
ash  and  as  a  mineral;  also  as  apatite. 

Phosphorus  is  prepared  by  heating  calcium  phosphate, 
silica  and  carbon  in  an  electric  furnace. 

Phosphorus  exists  as  white  and  red  phosphorus.  The 
former  is  very  poisonous. 

Tetraphosphorus  trisulfide  is  used  in  making  matches. 
Red  phosphorus  is  used  for  safety  matches. 


SUMMARY.    PHOSPHORUS,  ETC.  183 

Phosphine  is  prepared  by  warming  phosphorus  with  a 
solution  of  sodium  hydroxide.  So  prepared  it  .takes  fire 
in  the  air.  It  combines  with  hydriodic  acid. 

Phosphorus  pentoxide  is  prepared  by  burning  phosphorus 
in  dry  air.  It  is  an  effective  drying  agent. 

There  are  three  phosphoric  acids,  orthophosphoric, 
pyrophosphoric  and  metaphosphoric. 

Orthophosphoric  acid  is  a  tribasic  acid  but  only  the  first 
hydrogen  atom  ionizes  to  any  considerable  extent. 

Bimetallic  and  trimetallic  phosphates  are  hydrolyzed 
by  water  because  the  second  and  third  hydrogen  atoms 
have  only  faint  acid  properties.  Solutions  of  these  salts 
have  an  alkaline  reaction. 

Hydrolysis  is  the  decomposition  of  a  substance  by  water, 
especially  by  the  hydrogen  and  hydroxide  ions  of  the  water. 

Phosphorous  acid  is  dibasic  and  hypophosphorous  acid 
is  monobasic. 

Phosphorus  trichloride  and  phosphorus  pentachloricle  are 
hydrolyzed  to  phosphorous  and  phosphoric  acids  and 
hydrochloric  acid. 

Arsenic  is  found  in  copper  ores  and  as  mispickel,  arsenical 
pyrites. 

Arsenic  trioxide  is  used  as  a  ratsbane. 

Arsine  is  formed  when  an  arsenic  compound  is  introduced 
into  a  hydrogen  generator.  It  is  used  in  the  detection  of 
arsenic. 

Arsenic  trichloride  is  hydrolyzed  to  arsenious  oxide  and 
hydrochloric  acid. 

Arsenious  acid  is  known  only  in  very  dilute  solutions. 

Paris  green  is  a  double  salt  of  arsenious  acid  and  acetic 
acid. 

Arsenic  acid  is  tribasic.  The  silver  salt  and  the  magne- 
sium ammonium  salts  are  of  interest. 

Arsenic  trisulfide  is  lemon  yellow.  It  dissolves  in  am-^ 
monium  sulfide  to  give  ammonium  sulfarsenite. 


184  PHOSPHORUS,  ARSENIC,  ETC. 

Antimony  occurs  as  stibnite.  It  is  prepared  by  roasting 
the  ore  and  reducing  the  oxide  formed  with  carbon. 

Antimony  is  brittle  but  more  metallic  than  arsenic.  It 
is  used  in  type  metal  and  Babbitt  metal. 

Stibine  is  prepared  in  the  same  manner  as  arsine.  It 
decomposes  at  a  lower  temperature. 

Antimony  trichloride  is  hydrolyzed  to  antimony  oxy- 
chloride. 

Antimony  forms  three  oxides,  the  trioxide,  tetroxide 
and  pentoxide. 

Antimony  hydroxide  has  both  acid  and  basic  properties. 
Antimony  trisulfide  is  orange.  It  dissolves  in  ammonium 
sulfide  to  give  antimony  sulfantimonite. 

Bismuth  is  found  free.  It  is  used  in  stereotype  metal, 
fusible  metal  and  in  anti-friction  alloys. 

Bismuth  trioxide  does  not  dissolve  in  alkalies  and  the  tri- 
sulfide does  not  dissolve  in  ammonium  sulfide,  indicating 
that  bismuth  is  the  most  metallic  element  of  the  group. 

Bismuth  chloride  is  hydrolyzed  to  the  oxychloride  and 
bismuth  nitrate  to  a  basic  nitrate.  The  latter  is  used  in 
face  powders. 

EXERCISES 

1.  Prepare  a  table  of  the  hydrogen  compounds,  oxides,  chlorides, 
acids,  and  salts  of  the  sulfur  acids,  of  the  elements  of  Group  V. 

2.  Are  ammonia,  phosphine,  arsine  and  stibine  heavier  or  lighter 
than  air  and  in  approximately  what  proportions? 

3.  In  what  proportion  should  phosphine  and  oxygen  be  mixed 
for  an  explosive  mixture?     In  what  proportion  should  arsine  and 
oxygen  be  mixed? 

4.  Phosphorus    forms    an    oxychloride,    POC13.     How    many 
grams  of  phosphorus  pentachloride  and  how  many  grams  of  water 
would  be  required  to  prepare  100  grams  of  the  compound? 

5.  How  much  blue  vitriol,   CuS04.5H20,  will  be  required  to 
give  one  pound  (453  grams)  of  Paris  green? 


CHAPTER  XVII 
GROUP  IV:    CARBON 

Position  of  Carbon  among  the  Elements. — If  we  arrange 
the  elements  -of  low  atomic  weight  in  order,  omitting  hydro- 
gen, we  have  the  following  list : 

He,  4;  Li,  7;  Be,  9;B,  11;  C,  12;  N,  14;  O,  16;  F,  19;  Ne,  20 

In  this  list  carbon  stands  half  way  between  the  two  noble 
gases  helium  and  neon  and  also  half  way  between  the  strongly 
metallic  element  lithium  and  the  most  markedly  non-metallic 
element  fluorine.  In  its  properties  carbon  corresponds  to 
this  unique  position.  In  one  of  its  elementary  forms,  the 
diamond,  it  is  transparent  and  a  non-conductor  of  electricity . 
In  another  form,  graphite,  it  is  opaque,  has  a  metallic 
luster  and  is  a  fairly  good  conductor  of  electricity.  Car- 
bon also  shows  an  extraordinary  power  of  combining  with 
itself  and  as  a  result  of  this  power  it  forms  an  almost  end- 
less variety  of  complicated  compounds.  Many  of  these 
are  found  in  vegetable  and  animal  substances  and  carbon  is 
the  element  of  preeminent  importance  in  all  living  bodies. 

The  ultimate  source  of  all  of  the  carbon  for  the  growth  of 
trees  and  plants  is  the  carbon  dioxide  of  the  air.  Although 
it  forms  only  0.03  per  cent  of  the  volume  of  the  air,  growing 
plants  secure  their  supply  of  carbon  from  this,  absorbing 
from  sunlight  the  energy  by  means  of  which  they  seize  upon 
carbon  and  give  the  oxygen  back  to  the  atmosphere. 

Diamonds. — Diamonds  are  found  in  only  a  few  places  in 
the  world.  The  largest  supply  thus  far  discovered  is  at 
Kimberly,  in  South  Africa,  and  the  company  owning  the 

185 


186  CARBON 

mines  there  now  controls  the  diamond  markets  of  the 
world.  The  output  of  the  mines  is  valued  at  $15,000,000 
annually. 

The  diamond  is  the  hardest  substance  known.  It  also 
has  a  very  high  index  of  refraction  and  a  very  high  dispersive 
power;  that  is,  a  diamond  prism  separates  red  and  violet 
light  widely  in  the  spectrum.  Because  of  these  properties 
diamonds  which  are  cut  in  suitable  forms  reflect  and  re- 
fract the  light  which  falls  on  them  and  furnish  the  most 
brilliant  gems. 

Until  near  the  close  of  the  nineteenth  century  no  one  was 
able  to  discover  how  diamonds  were  formed  in  nature  or 
how  they  might  be  produced  artificially.  Finally,  the  French 
chemist  Moissan  discovered  some  microscopic  diamonds  in  a 
meteor  which  fell  in  Canon  Diablo,  Texas.  This  led  him  to 
the  thought  that  he  had,  as  it  were,  caught  nature  in  the 
act  of  making  diamonds.  After  considering  the  conditions 
under  which  the  diamonds  must  have  been  formed  in  the 
meteor  he  took  some  iron  and  melted  it  and  heated  it  to  a 
very  high  temperature,  above  the  melting  point,  in  an  elec- 
tric furnace.  He  then  dissolved  some  charcoal  in  the  iron 
and  plunged  the  mass  into  water.  This  caused  the  rapid 
solidification  of  the  exterior  while  the  interior  portions  were 
still  at  a  high  temperature  and  liquid.  Iron  expands  as  it 
solidifies  and  a  considerable  portion  of  the  dissolved  carbon 
separates,  ordinarily  in  the  form  of  graphite.  With  the  solid 
shell  on  the  outside  the  carbon  separated  from  the  iron 
under  conditions  of  high  pressure.  Moissan  demonstrated 
that  some  of  the  crystals  which  separated  were  diamonds, 
by  the  following  tests:  1.  The  crystals  sank  in  a  liquid  in 
which  graphite  will  float.  The  specific  gravity  of  the 
diamond  is  3.5;  that  of  graphite  is  2.25.  2.  The  crystals 
scratched  the  face  of  a  ruby.  Only  a  diamond  or  the  arti- 
ficial mineral  carborundum  will  do  this  and  the  specific 
gravity  of  the  crystals  was  greater  than  that  of  carborun- 


GRAPHITE  187 

dum.     3.  When  the  crystals  were  burned  in  a  current  of 
oxygen  12  parts  by  weight  gave  44  parts  of  carbon  dioxide. 

Diamonds  are  used  in  cutting  glass  but  an  edge  produced 
by  cleavage  must  be  employed.  Black  diamonds  of  an 
inferior  quality  are  used  for  drills.  These  are  used  espe- 
cially for  boring  in  such  a  manner  that  a  solid  core  of  rock 
may  be  removed.  By  this  means  a  solid  piece  of  rock  from 
the  bottom  of  a  deep  well  may  be  brought  to  the  surface 
and  examined. 

Graphite. — This  allotropic  variety  of  carbon  is  found 
much  more  abundantly  and  it  may  also  be  readily  prepared 
by  crystallizing  carbon  from  melted  cast  iron  and  by  heating 
impure  carbon  to  a  very  high  temperature  in  an  electric 
furnace.  Much  of  the  graphite  of  commerce  is  now  manu- 
factured in  this  manner. 

Graphite  is  sometimes  called  " black  lead"  and  "lead" 
pencils  are  made  from  it.  It  is,  of  course,  wholly  different 
from  metallic  lead.  The  graphite  for  lead  pencils  is 
specially  prepared  by  treatment  with  chemicals,  mixing1 
it  with  a  little  clay,  and  pressing  it  into  blocks  from  which 
the  leads  of  different  hardness  and  quality  are  to  be  cut. 

Graphite  is  used  as  a  lubricant,  especially  for  wooden 
bearings,  for  stove  polish  and  in  crucibles  designed  for  use 
at  very  high  temperatures  for  melting  steel  and  refractory 
alloys.  For  crucibles  the  graphite  is  mixed  with  a  small 
amount  of  clay  to  bind  the  particles  together.  The  graphite 
on  the  surface  of  the  crucible  burns  out  and  the  clay  then 
protects  the  remainder. 

Graphite  does  not  melt  even  at  the  temperature  of  the 
electric  arc  (about  3600°)  but  it  volatilizes  rapidly.  It 
also  sublimes  slowly  at  the  temperature  of  the  carbon  fila- 
ment in  an  electric  light  bulb. 

Amorphous  Carbon. — When  almost  any  compound  of 
carbon  or  almost  any  animal  or  vegetable  substance  is 
heated  to  a  high  temperature  it  turns  black  and  a  substance 


188  CARBON 

which  is  called  amorphous  carbon  separates.  It  is  only 
with  considerable  difficulty  that  by  heating  some  substance, 
such  as  sugar,  which  contains  only  carbon,  hydrogen  and 
oxygen,  a  pure,  amorphous  carbon  can  be  obtained.  Char- 
coal, lampblack,  coke  and  all  of  the  other  forms  of  amor- 
phous carbon  contain  hydrogen  and  are  otherwise  far  from 
pure  carbon. 

Amorphous  carbon  does  not  have  sharply  defined  prop- 
erties. The  density  varies  from  1.45  up  to  that  of  graphite. 
The  kindling  temperature  of  some  forms  is  as  low  as  300° 
while  that  of  others  approaches  that  of  graphite. 

Lampblack  is  prepared  by  burning  rosin,  petroleum  and 
similar  substances  under  conditions  to  produce  a  smoky 
flame.  It  is  used  as  a  pigment,  especially  in  printer's  ink. 
Its  value  depends  in  part  on  the  fact  that,  in  common  with 
other  forms  of  carbon,  there  is  nothing  that  will  dissolve  it 
at  ordinary  temperatures  and  it  is  absolutely  permanent 
and  unaffected  by  light. 

Charcoal. — When  wood  is  heated,  water,  tar  and  com- 
bustible gases  are  driven  out  and  there  finally  remains 
a  black,  porous  residue  retaining  the  form  of  the  wood 
and  consisting  chiefly  of  carbon  with  the  mineral  constit- 
uents originally  present  in  the  wood.  This  residue  is 
called  charcoal.  It  was  formerly  manufactured  by  piling 
wood  in  heaps,  covering  it  with  turf  and  allowing  it  to  burn 
with  a  smoldering  fire  until  the  conversion  to  charcoal 
was  complete.  This  wasteful  process  has  been  largely 
replaced  by  the  use  of  apparatus  for  heating  the  wood  in 
closed  retorts  in  such  a  manner  as  to  save  the  valuable 
products  which  distil  away,  as  well  as  the  charcoal.  Large 
quantities  of  charcoal  were  once  used  in  the  manufacture  of 
iron,  and  at  one  time  the  destruction  of  the  forests  in  England 
was  threatened  by  this  use.  It  has  now  been  displaced  by 
coal  and  coke.  It  is  still  sometimes  used  by  tinners  and 
it  is  used  for  the  filtration  of  alcohol  in  its  purification. 


COKE  AND  BY-PRODUCTS  J89 

Charcoal  will  condense  many  times  its  volume  of  am- 
monia, hydrogen  sulfide  and  other  gases,  especially  those 
gases  which  are  easily  liquefied.  The  gases  probably  con- 
dense to  the  liquid  form  in  the  minute  pores  of  the  charcoal. 
They  are  expelled  by  heat  and  freshly  heated  charcoal  is 
most  effective  for  the  absorption  of  the  gases. 

Different  kinds  of  charcoal  vary  greatly  in  their  absorp- 
tive power  and  especially  effective  kinds  made  from  coconut 
shells  have  been  developed  for  the  gas  masks  used  in  war. 

Bone-black  and  Animal  Charcoal  are  prepared  by  char- 
ring bones  and  slaughter-house  refuse.  They  are  used 
especially  to  remove  color  from  syrup  in  the  purification 
of  sugar. 

Coke  is  manufactured  by  heating  bituminous  coal. 
The  larger  part  of  the  coke  used  in  the  United  States  is 
made  in  beehive  coke  ovens,  round 
brick  chambers  about  12  feet  in 
diameter  and  7^  feet  high.  These 
are  partly  filled  with  the  coal  and 
the  gases  given  off  from  the  coal 
are  allowed  to  burn  above  the 
surface  of  the  coal  within  the 

chamber  (Fig.  34) .     The  burning  gases  furnish  the  heat  to 
coke  the  coal. 

By-products  Coke  Ovens. — In  Germany,  and  more  and 
more  in  this  country,  this  wasteful  method  is  being  replaced 
by  coking  ovens.  Flues  are  so  constructed  that  the  gas,  tar 
and  ammonia  water  may  be  recovered.  Nearly  all  of  the 
ammonia  of  commerce  comes  from  this  source.  The 
coal  tar  is  the  source  from  which  benzene,  toluene,  naph- 
thalene and  other  hydrocarbons  used  in  the  manufacture 
of  dyes  and  explosives  are  obtained.  Large  quantities  of 
the  tar  distillates  are  also  used  to  " creosote"  lumber  for 
railway  ties,  pavements  and  other  purposes  to  preserve 
it  from  decay.  The  gas  produced  is  more  than  enough  to 


190 


CARBON 


give  the  heat  required  for  coking  the  coal  and  may  be 
used  for  illuminating  gas  or  in  gas  engines  for  the  generation 
of  power.  Coke  is  chiefly  used  in  the  manufacture  of  iron. 
The  ovens  (Fig.  35),  are  so  arranged  that  the  air  for  the 
combustion  of  the  gas  used  in  heating  the  retorts  comes  up 
through  a  chamber  filled  with  a  checker- work  of  bricks. 
The  air  and  gas  come  together  and  burn  in  the  space 


FIG.  35. 

between  the  chambers  filled  with  coal.  The  heated  prod- 
ucts of  combustion  then  pass  down  through  a  second  set 
of  chambers  filled  with  bricks  which  absorb  the  heat  that 
would  otherwise  be  wasted.  After  a  time  the  current  of 
air  is  reversed,  going  up  through  the  heated  chamber  and 
down  through  the  cooled  one.  Such  an  arrangement  is 
called  a  " regenerative  furnace.'7 


GAS  CARBON.    COAL  191 

Gas  Carbon.  Carbon  Electrodes. — A  very  dense  form 
of  carbon,  approaching  graphite  in  its  properties,  is  pre- 
pared by  grinding  anthracite  coal,  petroleum  coke  and  other 
kinds  of  amorphous  carbon,  mixing  the  powder  with  a  little 
coal  tar  or  some  petroleum  product  to  make  it  cohere 
and,  after  subjecting  it  to  a  high  pressure,  heating  the  mix- 
ture to  1200°-1400°  for  one  or  two  days.  The  carbon  be- 
comes very  dense  and  hard  and  is  a  good  electrical  conductor. 
It  is  used  for  the  carbons  of  arc  lights  and  for  electrodes  in 
the  electrolytic  manufacture  of  chlorine  and  sodium  hydrox- 
ide from  brines. 

Coal. — During  millions  of  years  of  geological  time  vast 
stores  of  vegetable  materials  were  accumulated  in  different 
parts  of  the  world  from  the  growth  of  luxuriant,  tropical 
forests.  The  woody  fiber  of  this  material  contained  only 
about  50  per  cent  of  carbon  with  approximately  6  per 
cent  of  hydrogen  and  more  than  40  per  cent  of  oxygen. 
During  the  ages  which  have  passed  since  the  accumulation 
of  the  material  it  has  undergone  a  slow  process  of  carboniza- 
tion, by  which  the  oxygen  with  some  of  the  hydrogen  and  a 
little  carbon  have  escaped,  with  the  result  that  the  per  cent 
of  carbon  has  increased  while  the  per  cent  of  oxygen  has 
decreased,  but  the  per  cent  of  hydrogen  has  not  greatly 
changed  except  in  the  final  transformation  to  anthracite 
coal.  This  has  given  the  series  of  products  known  as  peat, 
lignite,  bituminous  coal  and  anthracite.  The  first  three 
give  off  volatile  matters  rich  in  hydrogen  and  carbon  when 
heated  and  for  this  reason  burn  with  a  smoky  flame  unless 
special  precautions  are  taken  to  secure  a  smokeless  combus- 
tion. Anthracite  coal  yields  almost  no  volatile  matters 
when  heated  and  burns  with  a  smokeless  flame. 

The  changes  which  have  occurred  in  these  transformations 
from  woody  fiber  to  anthracite  coal  can  be  seen  from  the 
following  table : 


192 


CARBON 


CHANGES  OF  WOOD  MATERIAL  DURING  GEOLOGICAL  TIME' 


| 

Percentage  composition 

i 

exclusive  of  moisture 

Calorific 

and  ash 

Percent-  i  Percent- 

value; 

Material 

age  of     1     age  of 

calories 

ash      '  i  moisture 

per  kilo- 

Car- 

Hydro- 

Oxy- 

Nitro- 

gram 

bon 

gen 

gen 

gen 

1 

|                    | 

Wood  —  Oak..  .  . 

50.35 

6.04 

43.52 

0.09 

0.37 

20.  OO2 

3696 

Peat  

59.70 

5.70 

33.04 

1.56 

11.84 

14.242 

3979 

Brown  Lignite, 

North  Dakota 

74.88 

4.99 

19.12 

1.01 

9.35      i      35.38 

3846 

Black    Lignite, 

| 

Colorado  

76.83 

5.34 

16.29 

1.54 

5.99      !      18.68 

5635 

Bituminous, 

Illinois.  ....... 

83.42 

5.29 

9.52 

1.77 

11.28      1        8.50 

6542 

Semibitumin- 

ous,  West  Vir- 

ginia       Poca- 

hontas  

91.50 

4.38 

3.07 

1.05 

6.55 

3.67 

7939 

Anthracite 

93.76 

2.72 

3.11 

0.41 

10.80 

2.18 

7216 

Charcoal  

84.11 

1.53 

14.36 

2.50 

6626 

Coke  

95.47 

0.67 

2.82 

1.04 

14.80 

6768 

1  Table  prepared  by  Professor  S.  W.  Parr.  2  Air  dry. 

There  are  three  kinds  of  bituminous  coals:  coking  coals, 
which  sinter  together  when  heated,  giving  a  hard,  coherent 
coke  suitable  for  use  in  blast  furnaces  for  the  manufacture 
of  cast  iron;  non-coking  coals,  which  do  not  sinter,  or  sinter 
imperfectly,  giving  a  friable  coke ;  and  cannel  coals,  having  a 
peculiar  homogeneous  structure  with  a  conchoidal  fracture. 
The  last  burn  with  a  brilliant,  luminous  flame  and  are  used 
in  the  manufacture  of  illuminating  gas. 

SUMMARY 

Carbon  is  at  the  center  of  the  first  period  of  elements  both 
in  position  and  properties.  It  is  the  most  important  element 
in  living  bodies. 

Plants  obtain  carbon  from  the  carbon  dioxide  of  the 
atmosphere. 

Diamonds  are  a  very  hard,  transparent,  dense  form  of  car- 
bon, valuable  for  gems,  cutting  glass  and  drills. 


EXERCISES;  CARBON  193 

Diamonds  have  been  prepared  artificially  by  crystallizing 
carbon  from  iron  under  a  high  pressure. 
,  Graphite  is  found  in  nature  and  is  prepared  by  crystalliz- 
ing carbon  from  a  solution  in  iron  or  by  heating  it  in  an 
electric  furnace. 

Graphite  is  used  in  lead  pencils,  stove  polish,  crucibles 
and  as  a  lubricant. 

Amorphous,  or  uncrystallized  carbon  is  found  in  an  im- 
pure form  in  lampblack,  charcoal,  coke  and  coal. 

Lampblack  is  used  as  a  pigment  and  in  printer's  ink. 

Charcoal  is  made  by  heating  wood.  It  is  used  for  small 
fires,  in  filtering  alcohol  and  to  absorb  noxious  gases. 

Bone-black  and  animal  charcoal  are  used  in  purifying 
sugar. 

Coke  is  made  by  heating  coal  in  beehive  or  in  by-product 
coke  ovens.  It  is  used  in  the  manufacture  of  iron. 

Gas  carbon  is  used  for  the  carbon  electrodes  of  arc  lights. 

Vegetable  matter  has  been  transformed  slowly  into  peat, 
lignite,  bituminous  coal  and  anthracite  coal,  and  in  some 
cases  to  graphite. 

Bituminous  coals  are  distinguished  as  coking,  non-coking 
and  cannel  coals. 

EXERCISES 

1.  The  heat  of  combustion  of  a  gram  atom  of  diamonds  is 
94,310  small  calories.     That  of  a  gram  atom  of  amorphous  carbon 
is  97,350  calories.     Is  the  formation  of  diamonds  from  amorphous 
carbon  exothermic  or  endothermic?     Will  the  formation  be  fav- 
ored by  an  increase  or  decrease  of  the  temperature  (see  p.  121)? 
Will  it  be  favored  by  pressure? 

2.  What  volume  of  carbon  dioxide  will  be  formed  by  burning  a 
gram  of  carbon?     What  volume  of  air  will  be  required  to  furnish 
the  oxygen  for  the  combustion? 

3.  What  volume  of  air  will  be  required  to  burn  a  pound  of 
carbon?     What  weight  of  air? 

13 


194  CARBON 

4.  What  volume  and  weight  of  air  are  required  to  burn  a  pound 
of  hydrogen? 

5.  What  volume  and  weight  of  air  will  be  required  to  burn  a 
pound  of  bituminous  coal  containing  73  per  cent  of  carbon  and 
4.7  per  cent  of  combustible  hydrogen? 

6.  How  many  kilograms  of  water  ought,  theoretically,  to  be 
evaporated  by  the  heat  from  burning  a  kilogram  of  coal  which 
gives  by  its  combustion  6500  calories  per  kilogram?     How  many 
pounds  of  water  per  pound  of  coal?     How  many  pounds  of  water 
evaporated  by  a  pound  of  coal  is  considered  good  boiler  efficiency? 


CHAPTER  XVIII 
HYDROCARBONS,  GAS,  FLAME 

Contrast  of  Carbon  with  Other  Elements. — The  elements 
of  the  halogen  family  each  form  only  a  single  compound 
with  hydrogen,  of  which  hydrochloric  acid,  HC1,  may  be 
taken  as  typical.  Oxygen  forms  two  compounds,  water, 
H2O,  and  hydrogen  peroxide,  H202.  Nitrogen  forms  four 
or  five  compounds  but  only  ammonia,  NH3,  is  met  with 
in  common  experience.  In  contrast  with  these,  carbon 
combines  with  hydrogen  to  form  several  thousands  of  com- 
pounds. No  other  element  gives  more  than  a  very  small 
number  of  compounds  with  hydrogen.  The  compounds  of 
hydrogen  with  carbon  are  called  hydrocarbons.  The  large 
number  of  these  is  due  to  the  power  which  carbon  has  of 
combining  with  itself  in  a  great  variety  of  ways. 

Series  of  Hydrocarbons. — The  compounds  of  carbon  and 
hydrogen  are  classified  by  arranging  them  in  a  series  in 
accordance  with  their  formulas  and  structure.  The  series 
which  contains  most  hydrogen  in  proportion  to  the  carbon  is 
called  the  marsh-gas  series  from  the  first  member,  methane 
or  marsh  gas,  CH4.  Some  members  of  the  series  are: 

Methane,  CH4 

Ethane,  C2H6 

Propane,  C3Hs 

Butane,  C4Hi0 

Pentane,  C5Hi2 

Hexane,  CeHu 

Heptane,  CrHie 

Octane,  CgHis 

Nonane,  CgH2o 

Decane,  CioH22 

195 


196  HYDROCARBONS,  GAS,  FLAME 

The  formula  CnH2n+2  may  be  used  for  any  member  of  the 
series,  n  standing  for  the  number  of  carbon  atoms. 

Structure  of  the  Hydrocarbons  of  the  Methane  Series.— 
It  will  be  noticed  that  in  the  successive  members  of  this 
series  each  new  carbon  atom  carries  with  it  two  new  hydro- 
gen atoms.  In  other  words,  there  is  a  difference  of  one 
carbon  atom  and  two  hydrogen  atoms  between  each  hydro- 
carbon of  the  series  and  the  one  preceding  or  following  it. 
The  same  relation  is  found  in  all  other  series  of  hydro- 
carbons. This  is  most  easily  explained  by  the  use  of  two 
well-established  principles:  1.  The  valence  of  carbon  is 
almost  always  four.  2.  Carbon  atoms  readily  unite  with 
each  other  as  well  as  with  other  elements. 

The  application  of  these  principles  leads  to  the  following 
structural  formulas  for  the  first  four  members  of  the  series : 

H  H      H  H      H      H 

II  III 

H-C-H       H-C-C-H      H-C-C-C-H 

I  11  111 

H  H      H  H       H      H 

Methane  Ethane  Propane 

H      H      H      H  H 

1         I         I        I  I 

H-C-C-C-C-H  H-C-H 


H 


H 


H      H      H      H 

Butane  TI         p          O          O         TT 

Boiling  point  +1°  '  ^   "  "   y   "  "  V   ' 

H      H      H 

Isobutane 
Boiling  point  —11.5° 

The  formulas  for  butane  and  isobutane  represent  two 
different  compounds,  C4Hi0,  which  are  actually  known. 
It  would  carry  us  too  far  to  explain  here  how  the  two  com- 
pounds are  distinguished. 

Methane,  CH4.  Fire-Damp.  Natural  Gas.— The  first 
member  of  the  series  is  often  called  marsh  gas.  If  the  decay- 


DA\  Y  SAFETY  LAMP 


197 


FIG.  36. 


ing  leaves  in  the  bottom  of  a  pond  are  stirred  up,  an  inflam- 
mable gas,  consisting  largely  of  methane,  usually  comes  up 
to  the  surface  and  may  be  collected  without  difficulty. 
The  gas  called  fire-damp,  which  often  causes 
disastrous  explosions  in  coal  mines,  consists 
largely  of  methane.  Natural  gas  also  in 
mostly  methane. 

Methane  is  most  easily  prepared  in  the 
laboratory  by  heating  a  mixture  of  sodium 
acetate  and  soda  lime.  The  latter  is  a 
mixture  of  sodium  hydroxide,  NaOH,  and 
slaked  lime,  Ca(OH)2. 

Davy  Safety  Lamp. — Mixtures  of  air  and 
methane  in  the  right  proportion  explode 
violently.  A  particularly  distressing  acci- 
dent in  a  coal  mine  from  this  cause,  early 
in  the  nineteenth  century,  led  to  a  request 
that  Sir  Humphrey  Davy  should  investigate  the  cause  of 
the  explosion  and,  if  possible,  suggest  a  remedy.  He 
established,  first  of  all,  that  when  marsh  gas  is  mixed  with 
from  six  to  fourteen  times  its  volume 
of  air  the  mixture  may  explode 
violently.  He  also  found  that  a  com- 
paratively high  temperature,  ap- 
proaching a  red  heat,  is  required  to 
ignite  the  mixture.  He  then  invented 
a  lamp  in  which  the  flame  is  com- 
pletely surrounded  with  heavy  wire 
gauze.  With  such  a  lamp,  Fig.  36, 
the  mixture  of  gases  on  the  outside  of 
the  gauze  does  not  become  heated  to 
its  kindling  temperature  and  so  does 
not  explode.  Precautions  are  always  taken,  however,  to 
avoid  the  presence  of  explosive  mixtures  in  mines.  Mix- 
tures of  coal  dust,  flour  dust,  cr  other  fine  particles  of  organic 


FIG.  37. 


198  HYDROCARBONS,  GAS,  FLAME 

matter,  with  air,  may  also  explode  violently.  An  explosion 
of  flour  dust  in  a  mill  in  Minneapolis  once  caused  the  de- 
struction of  the  mill  and  the  loss  of  several  lives.  Since 
then  great  care  is  taken  to  prevent  the  accumulation  of 
dust  in  flour  mills  and  factories.  The  effect  of  wire  gauze 
in  cooling  a  gas  below  its  kindling  temperature  is  shown 
in  Fig.  37. 

Petroleum. — Methane,  ethane,  propane  and  butane,  the 
first  four  members  of  the  methane  series,  are  gases  at  ordi- 
nary temperatures  but  the  higher  members  of  the  series  are 
liquids  or  solids.  A  number  of  these  are  found  in  petroleum, 
the  oil  which  is  obtained  in  many  places  by  boring  wells  in 
the  earth.  While  petroleum  has  been  known  for  a  very 
long  time  its  use  was  first  developed  on  a  large  commercial 
scale  in  Pennsylvania  in  1859.  Since  then  oil  has  been 
found  in  very  many  different  places  both  in  the  United 
States  and  in  foreign  countries.  Ohio,  Indiana,  Illinois, 
Texas,  Oklahoma,  Kansas  and  California  may  be  men- 
tioned as  states  producing  large  quantities  of  petroleum. 
Canada,  Mexico  and  the  region  of  the  Caucasus  may  be 
mentioned  among  foreign  countries.  There  is  reason  to 
think  that  immense  fields  of  petroleum  remain  undiscovered. 

Crude  petroleum  is  often  used  as  a  fuel.  It  is  purified 
for  the  production  of  kerosene,  gasoline  and  other  products, 
chiefly  by  distillation,  partly  by  treatment  with  concen- 
trated sulfuric  acid  to  remove  substances  which  give  the 
oi!  a  disagreeable  odor  or  interfere  with  its  proper  burning. 

Gasoline  consists  of  the  lower  boiling  constituents  which 
may  be  converted  into  a  vapor  at  a  sufficiently  low  tempera- 
ture to  form  an  explosive  mixture  with  air.  The  use  of 
gasoline  in  engines,  especially  in  automobiles,  is  familiar  to 
everyone.  The  successful  use  in  an  engine  depends  very 
largely  on  securing  a  proper  mixture  of  air  with  gasoline 
vapor.  If  too  much  gasoline  is  used,  a  part  burns  only  to 
carbon  monoxide  and  there  is  a  great  loss  of  energy.  Under 


UNSATURATED  COMPOUNDS 

some  conditions  a  part  of  the  carbon  fails  to  burn  at  all 
and  is  deposited  in  the  cylinders  of  the  engine.  The  carbon 
deposited  in  this  way  often  interferes  seriously  with  the 
operation  of  the  engine.  A  mixture  of  air  with  too  much 
gasoline  may  fail  to  explode  when  the  engine  is  hot,  just 
as  one  with  too  little  vapor  fails  to  explode  when  the  engine 
is  cold. 

Kerosene,  which  is  used  in  lamps  and  stoves,  should 
be  free  from  substances  which  are  volatile  enough  to  give 
an  explosive  mixture  with  air  at  moderate  temperatures. 
In  most  states  the  law  requires  that  the  flashing  point  of 
kerosene  shall  be  above  150°  F.,  i.e.,  so  much  vapor  that  it 
can  be  ignited  with  a  flame  must  not  be  given  off  below 
that  temperature.  With  the  increasing  demand  for  gaso- 
line, manufacturers  of  kerosene  are  no  longer  tempted  to 
sell  kerosene  with  a  low  flashing  point. 

Vaseline,1  paraffin  and  lubricating  oils  are  the  other  best 
known  products  obtained  from  petroleum. 

Ethylene,  C2H4.  Unsaturated  Compounds. — When  alco- 
hol, C2H6O,  is  heated  with  concentrated  sulfuric  acid  the 
elements  of  water  are  removed  and  a  gas,  ethylene,  C2H4,  is 
produced.  It  may  also  be  prepared  almost  quantitatively, 
by  passing  alcohol  over  aluminium  oxide  heated  to  a  mod- 
erate temperature. 

Ethylene  is  also  a  constituent  of  coal  gas  and  of  other 
illuminating  gases  and  is  one  of  the  most  important  of  the 
compounds  which  give  luminous  quality  to  such  gases. 

Ethylene  combines  directly  with  chlorine  or  bromine  to 
form  such  compounds  as  ethylene  chloride,  C2H4C12,  and 
ethylene  bromide,  C2H4Br2.  Because  of  this  property  it  is 
said  to  be  unsaturated.  Hydrocarbons  of  the  methane 
series,  which  do  not  act  in  the  same  way  but  which  give, 
instead,  substitution  products,  such  as  CH3C1  and  C2H5C1, 

1  Vaseline  is  a  proprietary  name  used  by  the  Chesebrough  Manufac- 
turing Company.  The  name  used  in  the  Pharmacopoeia  and  by  other 
manufacturers  is  petrolatum. 


200  HYDROCARBONS,  GAS,  FLAME 

are  said  to  be  saturated.     The  difference  will  be  clearer 
from  the  following  structural  formulas: 

H 

I  Cl 

H-C-H  | 

H-C  -H 
H-C-H  +  C1-C1=  | 

|  H-C-H  +  H-C1 

H  | 

Ethane  JJ 

Ethyl  chloride 

Cl 

H-C-H  H-C-H 

+  C1-C1=  | 

H-C-H  H-C-H 

Ethylene 

Cl 

Ethylene  dichloride 

Acetylene,  C2H2. — Calcium  carbide,  CaC2,  is  now  made 
on  a  large  scale  by  heating  a  mixture  of  lime  and  coke  in  an 
electric  furnace: 

CaO  +  3C  =  CaC2  +  CO 

Calcium      Carbon 
carbide         monoxide 

When  calcium  carbide  is  put  in  water  it  reacts  with  it, 
forming  acetylene,  C2H2,  and  calcium  hydroxide: 

CaC2  +  2HOH  =  Ca(OH)2  +  C2H2 

Acetylene  is  a  colorless  gas  which  burns,  under  proper 
conditions,  with  a  brilliant,  luminous  flame.  A  cubic  foot 
of  the  gas  may  be  made  to  give  more  than  ten  times  as  much 
light  as  a  cubic  foot  of  good  illuminating  gas. 

The  luminous  quality  of  the  acetylene  gas  flame  is  due 
to  the  separation  of  particles  of  carbon  which  are  heated  to 
a  very  high  temperature  in  the  flame.  A  special  burner, 
which  will  secure  complete  combustion  of  the  carbon  with- 
out the  escape  of  smoke,  is  required. 


ENDOTHERMIC  COMPOUNDS  201 

Various  forms  of  generators  for  acetylene  are  in  use.  The 
forms  in  which  the  carbide  is  dropped  into  water  are  more 
satisfactory  than  those  in  which  water  is  dropped  on  the 
carbide.  In  the  latter  forms  so  much  heat  is  generated 
that  a  part  of  the  acetylene  polymerizes  and  is  lost. 

Endothermic  Compounds. — When  acetylene  is  heated  it 
decomposes  into  carbon  and  hydrogen: 

O2ll2    =    2O    "f"    Il2 

Considerable  heat  is  evolved  during  the  decomposition. 
This  has  been  proved  by  comparing  the  heat  of  combustion 
of  acetylene  with  that  of  amorphous  carbon  and  hydrogen. 

26  g.  of  acetylene  give  when  burned 313.8  calories 

24  g.  of  carbon  give  when  burned  .  .195.3 
2  g .  of  hydrogen  give  when  burned  .  68 . 4     263 . 7 

Difference     50 . 1 

These  results  show  that  when  24  grams  of  carbon  combine 
with  2  grams  of  hydrogen  to  form  acetylene  50  calories  of 
heat  are  absorbed,  or  that  when  26  grams  of  acetylene 
decompose  into  carbon  and  hydrogen  50  calories  will  be 
given  off. 

These  facts  are  of  practical  importance  from  two  points 
of  view: 

1.  The  solid  carbon  which  separates  when  acetylene  is 
heated  is  the  source  of  the  intense  light  of  the  acetylene 
flame  (see  above). 

2.  If    the    decomposition    of    liquid    acetylene    is    once 
started,  the  heat  generated  by  the  decomposition  hastens 
it  and  the  large  volume  of  hydrogen  formed  may  cause  an 
explosion.     It  is,  in  fact,  possible  to  explode  liquid  acetylene 
with  a  cap,  very  much  as  nitroglycerine  is  exploded.     When 
these  properties  of  acetylene  became  generally  known  laws 
were  passed  forbidding  the  use  of  liquid  acetylene.     It  has 
been  discovered,  however,  that  acetone  or  acetaldehyde 


202  HYDROCARBONS,  GAS,  FLAME 

will  absorb  large  quantities  of  acetylene  under  pressure 
and  give  it  out  again  when  the  pressure  is  released  and  that 
the  solution,  if  not  too  rich  in  acetylene,  is  not  explosive. 
Such  a  solution  is  now  used,  especially  for  automobile  lights. 

Such  compounds  as  acetylene,  which  absorb  heat  in 
their  formation,  are  called  endothermic  compounds  and 
the  reactions  by  which  they  are  prepared  are  called  endo- 
thermic reactions. 

Reactions  of  this  type  are  favored  by  a  high  tempera- 
ture (p.  121),  and  in  accordance  with  this  acetylene  is 
formed  by  the  direct  union  of  carbon  and  hydrogen  at 
the  temperature  of  the  electric  arc. 

Benzene,  C6H6.  —  If  acetylene  is  passed  through  a  tube 
heated  to  a  moderate  temperature  a  part  of  it  will  poly- 
merize to  benzene: 


Benzene  and  several  other  closely  related  hydrocarbons 
are  found  in  coal  tar.  They  are  obtained  from  the  tar  by 
distillation.  These  compounds  are  used  in  large  quantities 
for  the  manufacture  of  dyes,  phenol  ("carbolic  acid") 
and  many  medicinal  products. 

Toluene,  C7H8,  the  second  compound  of  the  benzene 
series,  is  used  in  making  trinitrotoluene  ("T.  N.  T."), 
one  of  the  explosives  used  in  the  Great  War. 

Other  Hydrocarbons.  —  Each  of  the  hydrocarbons  men- 
tioned is  the  first  member  of  a  series.  Other  hydrocarbons 
of  the  ethylene  series  have  the  formulas  C3H6,  C4H8, 
C5Hi0,  C6Hi2,  etc.;  those  of  the  acetylene  series  are  C3H4, 
C4H6,  C5H8,  C6Hio,  etc.,  and  those  of  the  benzene  series 
are  C7H8,  C8Hi0,  C9Hi2,  Ci0Hi4,  etc.  Besides  these 
hydrocarbons  and  very  many  others  belonging  to  these 
four  series  there  are  many  other  series  and  several  thousand 
compounds  containing  only  carbon  and  hydrogen  are  known. 

Illuminating   Gas.  —  A  little  more  than  a  century  ago 


WATER  GAS  203 

the  only  forms  of  illumination  in  use  were  candles  and 
lamps  of  a  comparatively  inferior  type,  using  lard  or  some 
kind  of  vegetable  or  animal  oil.  Early  in  the  nineteenth 
century  the  manufacture  of  an  illuminating  gas  by  heating 
bituminous  or  cannel  coal  was  gradually  introduced  in 
large  cities.  The  gas  formed  in  this  way  is  a  complex 
mixture  of  hydrogen,  methane,  ethylene,  acetylene,  ben- 
zene vapor  and  other  hydrocarbons,  with  small  quantities 
of  carbon  monoxide,  carbon  dioxide  and  hydrogen  sulfide. 
The  hydrogen  sulfide  is  mostly  removed  before  the  gas 
is  used. 

When  gas  is  burned  in  a  flat  flame  or  in  an  Argand  burner 
the  illuminating  quality  depends  largely  on  the  so-called 
"  heavy  hydrocarbons,"  ethylene,  acetylene,  benzene,  etc., 
which  are  present.  If  burned  with  a  Welsbach  mantle 
the  light  given  is  nearly  proportional  to  the  heat  of  com- 
bustion of  the  gas,  but  the  heavy  hydrocarbons  give  much 
more  heat  in  proportion  to  their  volume  than  the  other 
constituents  do  and  are  still  important  in  determining  the 
quality  of  the  gas. 

Water  Gas. — For  the  manufacture  of  water  gas  a  suitable 
furnace  is  filled  with  a  large  mass  of  coke  and  this  is  brought 
to  a  high  temperature  by  burning  some  of  it  with  a  blast 
of  air.  The  air  is  then  cut  off  and  steam  is  turned  into 
the  chamber  for  a  few  minutes.  This  reacts  with  the  coke 
in  accordance  with  the  equation: 

C  +  H20  =  CO  +  H2 

This  reaction  is  an  endothermic  one  as  is  seen  from  the 
following  relations : 

12  grams  of  amorphous  carbon  give  97 . 65  calories 

28  grams  of  carbon  monoxide  give  68. 2 


2  grams  of  hydrogen  (burned  to 


steam)  give  58.0 

Excess  28 . 55  calorie; 


126.2 


204  HYDROCARBONS,  GAS,  FLAME 

In  order,  therefore,  to  convert  12  grams  of  carbon  and 
18  grams  of  steam  into  28  grams  of  carbon  monoxide  and 
2  grams  of  hydrogen,  28.5  calories  of  heat  must  be  added 
from  some  source.  Evidently  the  only  way  in  which  this 
heat  can  be  furnished  is  by  the  hot  coke  and  that  must  cool 
off  rapidly  as  the  reaction  proceeds.  After  a  few  minutes  it 
is  necessary  to  turn  off  the  steam  and  turn  on  the  blast  of 
air  to  bring  the  coke  again  to  a  high  temperature. 

While  the  steam  is  passing  in,  the  mixture  of  carbon 
monoxide  and  hydrogen  formed  is  saved  and  is  known  as 
water  gas.  Large  volumes  of  this  gas  can  be  manufactured 
very  cheaply  by  the  process. 

Pure  water  gas  burns  with  the  blue  flame  characteristic 
of  carbon  monoxide  and  gives  almost  no  light.  The  heat 
of  combustion,  too,  is  only  about  one-half  that  of  a  good 
illuminating  gas.  For  use  in  illumination  it  is  enriched  by 
adding  to  it  the  mixture  of  rich  gases  obtained  by  heating 
petroleum. 

The  most  serious  objection  to  the  use  of  water  gas  in 
illuminating  gas  is  the  poisonous  character  of  the  carbon 
monoxide  which  it  contains.  This  has  led  some  States  to 
prohibit  its  use.  There  is  additional  danger  from  the  fact 
that  water  gas  has  much  less  odor  than  coal  gas,  and  leaks 
are  not  so  quickly  noticed. 

Producer  Gas. — It  would  be  possible,  theoretically,  to 
pass  a  mixture  of  air  and  steam  over  a  mass  of  coke  or 
coal  in  such  proportions  that  the  burning  of  part  of  the 
coal  by  the  air  would  furnish  enough  heat  to  just  keep  up 
the  temperature  of  the  mass  and  still  convert  a  part  of 
the  steam  and  coal  into  water  gas.  The  gas  made  in  this 
manner  would  give  just  as  much  heat  in  its  combustion 
as  could  be  obtained  by  burning  the  original  coal.  Of 
course  some  heat,  must  be  lost  in  carrying  out  such  a 
scheme  practically,  but  apparatus  has  been  devised  by 
which  coal  may  be  converted  into  combustible  gases  that 


CANDLE  FLAME  205 

retain  80  to  85  per  cent  of  the  original  heating  power  of 
the  coal.  Such  a  gas  is  called  "  producer  gas."  It  consists 
chiefly  of  a  mixture  of  hydrogen,  carbon  monoxide,  CO, 
and  nitrogen. 

Producer  gas  can  be  used  much  more  economically  than 
the  original  solid  fuel  in  some  processes  for  the  manu- 
facture of  steel  and  glass,  and  in  gas  engines. 

Flames. — With  the  exception  of  the  mercury  vapor  elec- 
tric lamp,  all  forms  of  illumination  in  common  use  depend 
on  heating  some  solid  to  a  high  temperature. 

All  of  the  older  forms  of  illumination  depend  on  the 
combustion  of  gaseous  compounds  of  carbon  under  such 
conditions  that  solid  particles  of  carbon  exist  momentarily 
in  the  flame.  In  the  candle  or  lamp  the  material  burned 
is  converted  into  a  gas  immediately  before  combustion 
but  such  flames  are  gaseous  flames  just  as  much  as  a  flame 
of  illuminating  gas. 

Candle  Flame. — In  the  candle  flame  we  may  distinguish 
three  parts:  (1)  An  inner  portion  of  unburned  gas,  sur- 
rounding the  wick.  This  is  relatively  cool,  as  will  be  seen 
on  holding  a  wooden  toothpick  or  match  stick  across  the 
flame.  It  will  char  at  the  edges  of  the  flame  before  it  does 
in  the  center.  (2)  A  luminous  zone  of  partial  combustion. 
Here  the  carbon,  which  separates  in  the  decomposition 
of  compounds  present,  is  heated  red  hot  and  gives  the 
light  of  the  flame.  The  carbon  is  deposited  on  any  cold 
object  held  in  the  flame.  (3)  A  zone  of  complete  combus- 
tion surrounding  the  flame  and  most  apparent  at  the  base 
of  the  flame.  Here  the  carbon  and  hydrogen  of  the  flame 
are  completely  burned  to  carbon  dioxide  and  water. 

Bunsen  Burner. — In  the  Bunsen  burner,  Fig.  37  (p.  197), 
the  gas  is  mixed  with  an  amount  of  air  insufficient  for 
complete  combustion  before  it  is  ignited  at  the  top  of  the- 
burner.  This  causes  the  carbon  to  burn  to  carbon  monox- 
ide with  little  or  no  separation  of  carbon.  The  gases  burn 


206 


HYDROCARBONS,  GAS,  FLAME 


in  the  inner  cone  of  the  flame,  giving  a  mixture  of  carbon 
monoxide,  hydrogen,  carbon  dioxide,  water  vapor  and 
nitrogen.  In  the  outer  zone  of  the  flame  the  carbon  monox- 
ide and  hydrogen  burn  to  carbon  dioxide  and  water  vapor. 


27201 


Fio.  38. 


FIG.  39. 


Temperature  of  Flames, — The  temperature  of  the  flame 
'of  a  Bunsen  burner  varies  from  about  300°  in  the  center, 
near  the  mouth  of  the  burner,  where  combustion  has  not 
begun,  to  about  1550°  in  the  portion  between  the  inner  cone 
and  the  outside  of  the  flame.  These  temperatures  are 
shown  in  detail  in  Fig.  39. 


BLOWPIPE  207 

In  the  Meker  burner,  Fig.  38,  by  widening  the  top  of  the 
burner  and  giving  it  a  considerable  number  of  fairly  heavy 
metallic  partitions,  the  inner  cone  is  divided  into  a  number 
of  small  and  very  short  parts.  This  brings  the  high  tem- 
perature of  the  upper  part  of  the  Bunsen  flame  down  close 
to  the  mouth  of  the  burner,  concentrates  the  flame  and 
gives  it  a  more  uniform  and  somewhat  higher  temperature. 

The  temperatures  given  in  the  figures  are,  of  course,  the 
temperatures  of  the  flame  when  no  substance  radiating  heat 
is  present.  A  platinum  or  porcelain  crucible  placed  in  the 
flame  will  be  at  a  much  lower  temperature.  A  20-gram 
platinum  crucible  placed  1  cm.  above  the  Meker  burner, 
with  ordinary  gas,  will  usually  have  a  temperature  of- 
900°-950.° 

Blowpipe. — By  means  of  a  blowpipe,  Fig.  40,  the  flame  of 
a  candle,  or  the  luminous  flame  burning  at  the  slanting  end 
of  a  tube  introduced  in  a 
Bunsen  burner,  may  be 
changed  to  a  narrow, 
pointed  flame  which  can 
be  used  to  advantage  for 

either  the  oxidation  or  reduction  of  substances  laid  on  a 
stick  of  charcoal  or  dissolved  in  a  bead  of  borax  glass.  The 
extreme  tip  of  the  flame  gives  an  oxidizing  effect,  because  it  is 
intensely  hot  and  the  oxygen  of  the  air  can  act  on  the  sub- 
stance. The  inner,  faintly  luminous  cone  of  the  flame  is 
reducing  in  its  action,  owing  to  the  unburned,  combustible 
gases  present. 

SUMMARY 

Carbon,  in  contrast  with  other  elements,  forms  several 
thousands  of  compounds  with  hydrogen. 

Because  of  the  valences  of  carbon  and  hydrogen,  the 
successive  members  of  each  series  of  hydrocarbons  differ 
by  one  carbon  and  two  hydrogen  atoms. 


208  HYDROCARBONS,  GAS,  FLAME 

Methane  is  found  in  natural  gas  and  in  fire-damp.  It  is 
prepared  by  heating  sodium  acetate  with  soda  lime. 

Methane  and  other  hydrocarbons  give  explosive  mix- 
tures with  air.  Flour  dust  or  coal  dust  may  also  explode 
when  mixed  with  air  and  ignited. 

The  Davy  safety  lamp  was  invented  to  prevent  explosions 
in  coal  mines. 

Gasoline,  kerosene,  lubricating  oils,  vaseline  and  paraffin 
are  the  most  important  products  obtained  from  petroleum. 

The  flashing  point  of  gasoline  should  be  low,  for  satis- 
factory use,  that  of  kerosene  should  be  moderately  high, 
for  safety. 

Ethylene  is  prepared  by  the  decomposition  of  alcohol. 
It  gives  a  luminous  flame.  It  is  unsaturated  and  combines 
directly  with  chlorine  or  bromine. 

Calcium  carbide  is  prepared  by  heating  lime  and  carbon 
in  an  electric  furnace. 

Acetylene  is  prepared  by  the  action  of  water  on  calcium 
carbide.  It  is  an  endothermic  compound  and  decomposes 
into  its  elements  with  an  evolution  of  heat.  Liquid  acety- 
lene is  explosive. 

Benzene  and  many  other  hydrocarbons  are  obtained  from 
coal  tar.  These  are  used  in  making  dyes,  carbolic  acid  or 
phenol,  explosives  and  many  medicinal  products. 

Illuminating  gas  is  manufactured  from  coal  or  by  enrich- 
ing water  gas. 

Water  gas  is  made  by  passing  steam  over  hot  coke. 
The  reaction  is  endothermic  and  intermittent  for  that 
reason.  Water  gas  is  poisonous. 

Producer  gas  is  made  by  burning  coal  with  a  limited 
supply  of  air. 

The  luminosity  of  candle  flames  and  of  gas  flames  is  due 
to  heated  particles  of  carbon. 

In  the  Bunsen  burner  the  introduction  of  air  at  the  base 
of  the  burner  causes  the  partial  combustion  of  the  gas 


EXERCISES;  HYDROCARBONS 


209 


within    the    flame    and    this    prevents    the    separation    of 
carbon. 

A  blowpipe  may  be  used  to  give  either  an  oxidizing  or 
a  reducing  flame. 

EXERCISES 

1.  What  are  the  general  formulas  of  the  hydrocarbons  of  the 
ethylene,    acetylene   and    benzene    series    corresponding   to   the 
formula  CwH2n+2  for  the  methane  series? 

2.  What  is  the  general  formula  of  the  hydrocarbons  of  a  series 
intermediate  between  the  acetylene  and  benzene  series? 

3.  Write  the  equations  for  the  combustions  of  the  following 
hydrocarbons:  methane,   ethane,   acetylene,   benzene.     In   what 
proportion  by  volume  must  each  gas  be  mixed  with  oxygen  for 
complete  combustion?     In  what  proportion  with  air? 

4.  If  gasoline  has  an  average  composition  corresponding  to  the 
formula  C6Hi4,  in  what  proportion  by  volume  must  its  vapor  be 
mixed  with  air  for  complete  combustion? 

5.  What  weight  of  gasoline  will  be  required  to  form  the  most 
effective  explosive  mixture  with  one  cubic  foot  (28.3  liters)  of 
air? 

6.  Find  the  dimensions  of  the  cylinder  of  an  automobile  engine 
and  calculate  how  much  gasoline  should  be  introduced  for  one 
stroke  of  the  piston. 

7.  The  heats  of  combustion  of  one  cubic  foot  of  the  more  im- 
portant constituents  of  illuminating  gas  are  as  follows  in  calories 
and  in  British  Thermal  Units: 


Heat  of    t 

ombustion1 

Calories 

B.  T.  U. 

Carbon  monoxide,  CO. 

77    4 

307 

Hydrogen,  Ha  

66  3 

263 

Methane,  CH4  

215  0 

853 

Ethylene,  C2H4  

357  6 

1420 

Kiases  at  15.6°  (60°  F.)  burned  to  vapors  at  164°  (328°  F.). 
=3.968  British  Thermal  Units. 
14 


1  calorie 


210 


HYDROCARBONS,  GAS,  FLAME 


The  products  of  combustion  are  assumed  to  be  carbon  dioxide 
and  liquid  water. 

Assuming  that  the  "heavy  hydrocarbons"  consist  essentially 
of  ethylene,  what  will  be  the  heat  of  combustion  of  one  cubic  foot 
of  the  following  samples  of  gas: 


Coal  gas 

Enriched 
water  gas 

Producer  gas 

Carbon  dioxide,  C02. 

1    1 

3  0 

1  5 

Carbon  monoxide,  CO  . 

7  2 

26  1 

23  5 

Hydrogen,  HS 

49  0 

32  1 

6  0 

Methane,  CH4  

34.5 

19.8 

3.0 

Heavy  hydrocarbons. 

5  0 

16  6 

Nitrogen  

3.2 

2.4 

66.0 

Candle-power  

100.0 
17.5 

100.0 
25.0 

100.0 

8.  The  heat  of  combustion  of  one  gram  atom  of  carbon  is 
97,650  calories  (small).  The  heat  of  combustion  of  one  gram 
molecule  of  hydrogen  burned  to  liquid  water  is  68,400  calories. 
The  heat  of  combustion  of  one  gram  molecule  of  methane,  CH4, 
burned  to  carbon  dioxide  and  liquid  water  is  214,000  calories.  Is 
the  reaction  expressed  by  the  equation: 

C  +  2H2  «=*  CH4 

exothermic   or   endothermic?     Will   the   equilibrium   be   shifted 
toward  the  formation  of  methane  by  a  low  or  a  high  temperature? 


CHAPTER  XIX 

CARBON    MONOXIDE,    CARBON    DIOXIDE,    CARBON 
DISULFIDE,  CYANIDES 

Formation  of  Carbon  Monoxide. — Whenever  fuels  burn 
in  a  thick  layer  and  especially  in  the  burning  of  anthracite 
coal,  blue  flames  will  be  seen  playing  over  the  surface. 
These  are  flames  of  carbon  monoxide  burning  to  carbon 
dioxide.  In  the  lower  part  of  the  mass  of  fuel  some  of  it 
may  be  burned  to  carbon  dioxide  and  the  latter  is  reduced 
to  the  monoxide  as  it  passes  through  the  hot  coals  above: 

CO2  +  C  =  2CO 

The  reaction  is  an  endothermic  one  and  only  occurs  at  a 
high  temperature.  Twelve  grams  of  carbon  burned  to 
carbon  monoxide  give  29.45  calories  (large),  while  the  same 
weight  of  carbon  burned  to  carbon  dioxide  gives  97.65 
calories.  From  this  relation  it  will  be  seen  that  conditions 
of  combustion  in  a  furnace  or  in  a  gas  engine  which  lead  to 
the  escape  of  carbon  monoxide  are  very  wasteful  of  the 
energy  of  the  fuel.  It  is  partly  for  the  same  reason  that  it 
is  possible  to  convert  coal  into  a  producer  gas  which  retains 
a  large  per  cent  of  its  original  energy  (p.  204). 

Preparation  and  Properties  of  Carbon  Monoxide. — 
Carbon  monoxide  is  most  easily  prepared  in  the  laboratory 
by  heating  a  mixture  of  oxalic  acid  and  concentrated  sul- 
furic  acid: 

H2C2O4  +  H2S04  =  CO  +  C02  +  H2SO4.H2O 

Oxalic  acid 

The  sulfuric  acid  catalyzes  the  decomposition  of  the 
oxalic  acid  and  combines  with  the  water  formed.  The 

211 


212  CARBON  OXIDES,  CYANIDES 

carbon  dioxide  formed  at  the  same  time  may  be  removed  by 
passing  the  gas  through  a  tube  filled  with  soda  lime. 

Carbon  monoxide  is  a  colorless,  odorless  gas.  It  burns 
with  a  characteristic  blue  flame  that  gives  very  little 
light. 

Carbon  monoxide  is  a  very  dangerous  poison.  Air  con- 
taining one  part  in  a  thousand  of  the  gas  is  unsafe  to  breathe, 
even  for  a  short  time.  Air  containing  smaller  amounts 
may  cause  serious  injury  in  the  air  of  living  rooms.  The  gas 
seems  to  combine  with  the  hemoglobin  of  the  blood  and  to 
so  change  it  that  it  is  no  longer  capable  of  performing  its 
normal  function  of  transferring  the  oxygen  of  the  air  to  the 
tissues  of  the  body.  For  this  reason  it  acts  as  a  cumulative 
poison  and  recovery  from  poisoning  with  the  gas  is  very 
slow.  This  property  of  carbon  monoxide  constitutes  a  very 
serious  objection  to  the  use  of  water  gas. 

The  escape  of  carbon  monoxide  from  a  charcoal  fire  and 
from  a  flame  burning  against  a  cold  surface  is  the  reason 
why  the  products  of  combustion  of  neither  should  be  per- 
mitted to  escape  into  a  living  room. 

Carbon  Dioxide. — The  formation  of  carbon  dioxide  by  the 
burning  of  carbon  and  its  compounds  and  by  respiration 
have  been  repeatedly  referred  to.  The  gas  is  most  easily 
prepared  by  the  treatment  of  a  carbonate  with  hydrochloric 
or  sulfuric  acid.  Carbonic  acid,  H2CO3,  very  readily  de- 
composes into  its  anhydride,  CO2,  and  water,  and  as  carbon 
dioxide  is  not  very  soluble  in  water  the  escape  of  the  carbon 
dioxide  causes  the  reactions  to  go  nearly  to  completion  in 
the  direction  of  its  formation : 

CaCO3  +  2HC1  <=±  CaCl2  4-  H2CO3 

Calcium 
carbonate 

H2CO3  <=±  CO2  +  H2O 

Carbon  dioxide  is  a  colorless  gas  about  one-half  heavier 
than  air.  It  is  not  distinctly  poisonous  when  present  in 


HENRY'S  LAW  213 

moderate  amounts.  It  may,  however,  cause  death  from 
suffocation.  The  gas  sometimes  accumulates  in  wells  and 
mines  and  is  known  as  choke-damp.  "Where  its  presence 
is  suspected  the  place  should  be  tested  with  a  burning 
candle  before  it  is  entered,  though  it  is  possible  to  live  for  a 
short  time  in  a  room  containing  so  much  carbon  dioxide  that 
a  candle  is  extinguished. 

Carbon  Dioxide  and  Water.  Henry's  Law.  Soda  Water. 
— -At  ordinary  temperature  and  atmospheric  pressure, 
water  dissolves  about  its  own  volume  of  carbon  dioxide. 
From  air  containing  carbon  dioxide  the  gas  will  be  dissolved 
in  proportion  to  the  " partial  pressure"  of  the  gas  present. 
Thus  the  "partial  pressure"  of  carbon  dioxide  in  air  con- 
taining 10  per  cent  of  the  gas  by  volume  is  one-tenth  of  an 
atmosphere  and  only  one-tenth  as  much  carbon  dioxide  will 
be  absorbed  from  such  a  mixture  as  will  be  absorbed  if 
water  were  in  contact  with  the  pure  gas.  This  is  known  as 
Henry's  Law  and  applies  to  all  gases  that  are  slightly 
soluble  in  water. 

Under  pressure  water  still  takes  up  its  own  volume  of 
the  gas  but  this  means,  of  course,  that  the  weight  of  the 
gas  absorbed  increases  proportionally  with  the  pressure. 
When  the  pressure  is  removed  the  excess  of  gas  escapes  with 
effervescence.  This  property  is  used  in  soda  water  and  in 
carbonated  beverages. 

The  solution  of  carbon  dioxide  in  water  has  a  faintly 
acid  taste  and  reddens  blue  litmus.  The  taste  is  only 
slightly  acid  because  the  ionization : 

-   H2CO3  <=*  H+  +  HCO3- 

is  only  very  slight,  most  of  the  carbon  dioxide  remaining 
either  as  carbonic  acid,H2C03,or,  perhaps,  as  carbon  dioxide 
uncombined  with  the  water.  Solutions  of  sodium  carbonate, 
Na2COa,  and  potassium  carbonate,  E^COs,  have  an  alkaline 


214  CARBON  OXIDES,  CYANIDES 

reaction.  Explain  this  as  has  been  done  for  alkaline  phos- 
phates. 

Solutions  of  carbon  dioxide  in  water  react  with  bases  to 
form  carbonates.  The  effect  of  the  gas  on  lime  water, 
Ca(OH)2,  has  been  referred  to  several  times  (p.  11). 

Carbon  dioxide  can  be  liquefied  by  pressure  and  the 
liquid  is  sold  in  strong  steel  cylinders  for  use  in  soda-water 
fountains  and  for  carbonating  beverages.  The  pressure 
in  these  cylinders  is  from  60  to  70  atmospheres,  or  900  to 
1000  pounds  to  the  square  inch.  If  the  liquid  is  allowed  to 
escape  into  a  strong  cloth  bag  a  part  of  it  evaporates  at 
once  as  the  pressure  is  removed  and  the  remainder  is  cooled 
to  so  low  a  temperature  that  it  freezes  to  a  white,  snow- 
like  solid.  The  temperature  of  this  solid  is  —  79°,  the  boiling 
point,  or  rather  subliming  point,  of  solid  carbon  dioxide. 
In  other  words,  the  vapor  pressure  of  solid  carbon  dioxide  is 
760  mm.  at  -  79°.  The  melting  point  of  the  solid  is  -  56.4°, 
more  than  20°  above  the  boiling  point.  The  vapor  pres- 
sure at  the  melting  point  is  5.1  atmospheres. 

Solid  carbon  dioxide  is  a  very  convenient  means  of 
securing  low  temperatures  in  the  laboratory. 

Carbon  Bisulfide. — When  the  vapor  of  sulfur  is  passed 
over  heated  charcoal  the  sulfur  combines  with  the  carbon 
to  form  carbon  disulfide,  CS2.  The  preparation  is  now 
carried  out  in  an  electric  furnace. 

Carbon  disulfide,  when  pure,  is  a  colorless  liquid  which 
boils  at  47°  and  gives  off  a  considerable  amount  of  vapor 
at  ordinary  temperatures.  The  kindling  temperature  is 
very  low,  and  mixtures  of  the  vapor  with  air  are  highly 
explosive.  For  this  reason  carbon  disulfide  must  be  handled 
with  extreme  care.  This  inflammability  interferes  with 
its  use  as  a  solvent  for  fats  and  for  some  purposes  to  which 
it  is  otherwise  well  adapted.  It  is  used  in  vulcanizing 
india-rubber,  in  rubber  cements  and  some  times  as  a  poison 
to  kill  rats,  ground  squirrels  and  moths. 


CYANIDES  215 

Potassium  Ferrocyanide. — When  a  mixture  of  slaughter 
house  refuse,  rich  in  nitrogenous  carbon  compounds,  is 
heated  with  potassium  carbonate,  K2CO3,  and  iron  turn- 
ings the  elements  unite  to  form  potassium  ferrocyanide, 
K4Fe(CN)6.  This  is  a  compound  formed  by  the  union  of 
potassium  cyanide,  KCN,  and  ferrous  cyanide,  Fe(CN)2, 
but  it  is  also  to  be  considered  as  a  salt  of  the  acid, 
H4Fe (CN)  6,  called  hydroferrocyanic  acid.  Potassium  ferro- 
cyanide dissolves  readily  in  water  and  crystallizes  from 
the  solution  as  a  yellow  hydrate,  having  the  composition, 
K4FeC6N6.H2O. 

Cyanides. — When  potassium  ferrocyanide  is  heated  with 
sodium  the  iron  is  replaced  by  the  sodium,  giving  a  mixture 
of  potassium  cyanide,  KCN,  and  sodium  cyanide,  NaCN: 

K4FeC6N6  +  2Na  =  4KCN  +  2NaCN  +  Fe 

A  solution  of  these  cyanides  will  dissolve  metallic  gold, 
if  used  in  the  presence  of  air,  and  such  a  solution  is  exten- 
sively used  in  extracting  gold  from  its  ores. 

Hydrocyanic  Acid  or  Prussic  Acid. — If  a  solution  of 
potassium  ferrocyanide,  K4FeC6N6  or  of  potassium  cyanide, 
KCN,  is  mixed  with  dilute  sulfuric  acid  and  distilled, 
hydrocyanic  acid,  which  is  volatile,  will  pass  over.  Pure 
hydrocyanic  acid  is  a  volatile  liquid  which  boils  at  26.5°. 
It  is  one  to  the  quickest,  most  powerful  poisons  known. 
The  cyanides  are  also  very  poisonous. 

A  dilute  solution  of  hydrocyanic  acid  is  sometimes  used 
in  medicine. 

Complex  Cyanides. — Many  of  the  cyanides  of  heavy 
metals  are  insoluble  in  water  but  most  of  these  cyanides 
will  dissolve  in  a  solution  of  potassium  cyanide.  In  the 
resulting  solution  the  heavy  metal  enters  into  a  complex 
group  which  reacts  as  a  whole  and  no  longer  shows  the 
characteristics  of  the  heavy  metal  which  it  contains.  Thus 
potassium  ferrocyanide,  K4FeC6Ne,  will  give  no  precipitate, 


216  CARBON  OXIDES,  CYANIDES 

with  sodium  hydroxide,  NaOH,  ammonium  sulfide,  (NH4)2S, 
or  with  other  reagents  which  react  readily  with  ferrous 
sulfate,  FeSC>4,  or  with  ordinary  ferrous  salts. 

If  an  electric  current  is  passed  through  a  solution  of  ferrous 
sulfate,  FeSO4,  the  ferrous  ion,  Fe++,  is  carried  toward 
the  negative  pole,  or  cathode,  and  the  sulfate  ion,  SO4=, 
is  carried  toward  the  positive  pole,  or  anode.  If  the  current 
is  passed  through  a  solution  of  potassium  ferrocyanide, 
however,  the  iron  is  not  carried  toward  the  cathode  but 
it  is  carried,  instead,  with  the  cyanogen,  CN,  toward  the 
anode  and  only  the  potassium  travels  toward  the  cathode. 
This  is  best  explained  by  supposing  that  the  ions  in  the 
solution  are  not  4K+,  Fe++  and  6CN~,  as  might  have  been 
expected,  but  that  they  are  4K+  and  FeC6N6E.  In  other 
words  the  iron  and  cyanogen  unite  to  form  a  complex  ion 
which  is  called  the  ferrocyanide  ion. 

Many  other  complex  cyanides  are  known.  Among  these 
may  be  mentioned  silver  argenticyanide,  KAgC2N2  (or 
KCN.AgCN),  which  is  used  in  silver  plating,  and  potas- 
sium ferricyanide,  K3FeC6N6,  or  3KCN.Fe(CN)3,  a  red 
salt  prepared  by  oxidizing  potassium  ferrocyanide.  The 
last  salt  is  used  in  preparing  blue-print  paper. 

SUMMARY 

Carbon  monoxide  is  formed  when  carbon  dioxide  passes 
over  hot  carbon.  The  reaction  is  endothermic. 

Oxalic  acid  when  heated  with  concentrated  sulfuric 
acid  gives  carbon  monoxide  and  carbon  dioxide.  The 
carbon  dioxide  may  be  absorbed  by  soda  lime. 

Carbon  monoxide  burns  with  a  blue  flame.  It  is  very 
poisonous.  It  is  a  constituent  of  water  gas. 

Carbon  dioxide  is  formed  by  burning  carbon.  It  is 
prepared  by  the  action  of  an  acid  on  a  carbonate. 

Water  dissolves  about  its  own  volume  of  carbon  dioxide 


EXERCISES.     CARBON  OXIDES  217 

whether  under  high  or  low  pressure  (Henry's  Law)  and  so 
dissolves  a  greater  weight  under  high  pressures.  The 
solution  contains  carbonic  acid,  which  is  a  very  weak  acid. 
The  salts  of  the  alkali  metals  have  an  alkaline  reaction  in 
solution. 

The  subliming  point  of  pure  carbon  dioxide  is  —79°. 
The  melting  point  is  more  than  20°  higher.  Liquid  carbon 
dioxide  is  possible  only  under  pressure. 

Carbon  disulfide  is  prepared  by  passing  sulfur  vapor  over 
hot  charcoal.  It  is  poisonous  and  very  inflammable.  It 
is  used  in  vulcanizing  rubber,  and  sometimes  as  a  poison. 

Potassium  ferrocyanide  is  prepared  by  heating  nitro- 
genous matter  with  potassium  carbonate  and  iron. 

A  mixture  of  potassium  and  sodium  cyanides  is  prepared 
by  heating  potassium  ferrocyanide  with  sodium.  It  is  used 
in  extracting  gold  from  its  ores. 

Hydrocyanic  or  prussic  acid  and  the  cyanides  are  very 
poisonous. 

Potassium  cyanide  combines  with  cyanides  of  the  heavy 
metals  to  form  complex  cyanides  in  which  the  heavy  metal 
forms  a  part  of  the  anion.  Solutions  of  such  complex 
cyanides  often  fail  to  give  the  ordinary  reactions  used  to 
detect  the  heavy  metals  which  they  contain. 

Potassium  argenticyanide  is  used  in  silver  plating. 

EXERCISES 

1.  Write  the  equation  for  the  action  of  dilute  sulfuric  acid  on 
potassium  ferrocyanide. 

2.  When  a  slightly  diluted  sulfuric  acid  is  heated  with  potassium 
ferricyanide,  carbon  monoxide  is  formed.     The  other  products  are 
ammonium  sulfate,  potassium  sulfate  and  ferric  sulfate.     Write 
the   equation.     Notice   that   two  molecules  of  the  ferricyanide 
must  be  used  (why?)  and  enough  water  to  furnish  the  oxygen  of  the 
carbon  monoxide. 

3.  In   what   proportion  by  volume  should   carbon  monoxide 
and  oxygen  be  mixed  for  explosion?     When  cold  what  will  be  the 


218  CARBON  OXIDES,  CYANIDES 

volume  of  the  gaseous  product  as  compared  with  volume  of  the 
original  gases? 

4.  If  the  mixture  of  carbon  monoxide  and  oxygen  gives  a  tem- 
perature of  2500°,  what  will  be  the  volume  of  the  carbon  dioxide 
formed  at  that  temperature  as  compared  with  the  volume  of  the 
mixed  gases  at  20°  before  the  explosion? 

5.  Sodium  carbonate,  Na2C03,  is  hydrolyzed  by  water,  in  part, 
giving  hydrocarbonate,  HC03~  and  hydroxide,  OH~,  ions.     Write 
the  equation  for  the  reaction-  between  the  ions  of  water  and  the 
ions  of  sodium  carbonate  showing  the  ions  which  result.     What 
will  be  the  reaction  of  the  solution?     Confirm  this  by  testing  a 
solution  of  sodium  carbonate  with  litmus  paper. 

6.  What  per  cent  of  the  heat  energy  of  carbon  is  lost  when  it  is 
burned  only  to  carbon  monoxide? 


CHAPTER  XX 

CARBOHYDRATES,    ALCOHOLS,    ACIDS,    BREAD,   PRO- 
TEINS, DIGESTION,  ANTITOXINS,  ALKALOIDS,  DYES 

Carbohydrates. — Among  the  many  thousands  of  com- 
pounds which  contain  carbon,  hydrogen  and  oxygen,  those 
of  one  of  the  most  important  groups  are  called  carbohydrates 
because  the  hydrogen  and  oxygen  in  them  are  in  the  same 
proportion  as  in  water.  The  name  seems  to  imply  that 
they  are  hydrates  of  carbon,  that  is,  that  they  are  formed 
by  the  union  of  carbon  with  water.  Such  a  point  of  view 
is  not  justified,  for  they  are  not  formed  in  nature,  or  arti- 
ficially, by  the  direct  union  of  carbon  with  water,  and  when 
heated,  while  carbon  and  water  are  formed  by  their  decom- 
position, many  other  substances  are  formed  as  well. 

The  carbohydrates  include,  especially,  cellulose,  starch 
and  many  different  sugars. 

Cellulose,  is  represented  by  the  formula  (C6HioO5)n. 
No  means  has  been  discovered  for  determining  the  value  of 
n  in  this  formula,  because  cellulose  cannot  be  vaporized 
without  decomposition  and  there  is  no  simple  solvent  in 
which  it  might  be  dissolved  and  its  molecular  weight  deter- 
mined by  the  lowering  of  the  freezing  point  or  rise  of  the 
boiling  point  of  its  solution  as  is  done  with  many  other 
compounds  which  cannot  be  vaporized. 

Cellulose  forms  the  larger  part  of  the  woody  fiber  of  trees 
and  all  kinds  of  plants. 

In  the  form  of  grass,  clover,  alfalfa  and  the  hay  or  silage 
made  from  these  or  from  corn  it  is  an  important  food  for 
herbivorous  animals.  It  is  a  constituent  of  many  foods 

219 


220  CARBOHYDRATES,  ALCOHOLS,  ETC. 

used  by  man  but  apparently  it  is  not  digested  and  utilized 
to  any  appreciable  extent. 

Coal  was  formed  largely  from  the  woody  fiber  of  plants 
which  grew  ages  ago.  Coal  and  wood  form,  of  course,  our 
most  important  fuels.  There  is  some  probability  that 
natural  gas  and  petroleum  were  formed,  in  part,  from 
cellulose. 

Cotton,  linen,  hemp  and  other  vegetable  fibers  used  in 
the  manufacture  of  cloth,  ropes  and  twine  are  largely  com- 
posed of  cellulose. 

Paper  is  almost  entirely  cellulose.  The  cheaper  grades, 
used  for  printing  the  daily  papers,  are  made  from  wood. 
Better  kinds  of  paper  are  made  from  linen  rags  and  other 
fibrous  materials.  The  materials  are  bleached  and  purified 
by  the  use  of  various  chemicals  and  mixed  with  water  to  a 
thin  pulp  which  can  be  spread  out  in  a  uniform  layer,  which 
is  then  pressed  and  dried. 

Nitrocellulose;  Gun  Cotton;  Lacquers;  Collodion;  Arti- 
ficial Silk. — When  cotton,  which  is  nearly  pure  cellulose, 
is  treated  with  a  mixture  of  concentrated  sulfuric  and 
nitric  acids  it  is  converted  into  a  mixture  of  compounds 
called  usually  nitrocellulose,  but  more  correctly  cellulose 
nitrate.  These  are  formed  by  the  replacement  of  hydroxyl 
groups,  OH,  by  the  nitrate  group,  NO3,  just  as  sodium 
nitrate  is  formed  by  the  replacement  of  the  hydroxyl  group 
of  sodium  hydroxide: 

NaOH  +  HNO3  =  NaN03  +  HOH 
Ci2H14O4(OH)6  +  6HNO3  =  C12H14O4(NO3)6  +  6HOH 

Cellulose  hexanitrate 

Cellulose  hexanitrate  is  the  powerful  explosive  known  as 
gun-cotton  and  also  used  as  the  basis  of  the  smokeless 
powders. 

Other  nitrates  are  formed  by  the  replacement  of  a  smaller 
number  of  hydroxyl  groups.  Solutions  of  some  of  these  in 


CELLULOID.     STARCH 


221 


amyl  acetate  or  other  solvents  are  excellent  lacquers  for 
brass.  Collodion  is  a  solution  in  ether  and  alcohol.  Arti- 
ficial silk  is  made  from  cellulose  nitrate  or  acetate  but  if  the 
former  is  used  it  is  subjected  to  a  treatment  which  removes 
the  nitrate  group  and  makes  it  less  inflammable. 

Celluloid  is  a  mixture  of  some  of  the  cellulose  nitrates 
with  camphor.  It  is  highly  inflammable  but  not  explosive 
in  the  ordinary  sense. 


$fr$ 


o^x^.R  OWQf* 


FIG.  41. — A,  potato  starch;  B,  rice  starch;  C,  wheat  starch 
(X  160).     After  Allen. 

Starch  also  has  the  formula  (C6H10O5)n  and  its  molecular 
weight  is  unknown  It  is  found  in  the  form  of  granules  of 
various  sizes  and  shapes  in  wheat,  maize,  rice,  sago,  tapioca 
and  many  other  substances  used  as  articles  of  food.  It  is 
the  most  important  non-nitrogenous  compound  in  articles 
of  human  diet  and  it  furnishes  a  considerable  portion  of  the 
heat  of  our  bodies  and  of  the  muscular  energy  with  which  we 
move  and  do  work. 


'222  CARBOHYDRATES,  ALCOHOLS,  ETC. 

The  forms  of  the  granules,  some  of  which  are  illustrated 
in  Fig.  41  have  no  connection  with  the  chemical  composition, 
which  seems  to  be  the  same  in  all  plants.  A  very  thin,  outer 
shell  covering  the  granules  is  probably  of  a.  somewhat  dif- 
ferent character  from  the  starch,  but  it  has  not  been  found 
possible  to  separate  and  examine  it. 

When  foods  containing  starch  are  cooked  the  cell  walls 
are  burst  and  the  starch  forms  a  soft,  pulpy  mass,  which  is 
easily  attacked  by  the  digestion  fluids.  Pure  starch  forms 
a  paste  with  hot  water.  With  larger  amounts  of  water  it 
gives  a  slightly  opalescent,  colloidal  solution. 

Starch  gives  an  intense  blue  color  with  a  solution  of  iodine 
in  potassium  iodide.  This  is  used  as  a  sensitive  test  for 
starch  or  for  iodine. 

Sucrose  or  Cane  Sugar.  Beet  Sugar. — The  juices  of 
isugar  cane,  beets,  the  sap  of  maple  trees  and  nearly  all 
fruits  contain  a  crystalline,  easily  soluble  compound  hav- 
ing the  composition  C^H^On.  It  is  commonly  known 
merely  as  sugar,  more  accurately  as  sucrose,  or  cane 
sugar;  It  is  manufactured  chiefly  from  sugar  cane  and 
sugar  beets. 

The  juice  of  the  sugar  cane  is  pressed  out  with  powerful 
rolls.  It  is  then  concentrated  to  the  point  of  crystalliza- 
tion by  evaporation  under  diminished  pressure.  The  reduc- 
tion of  the  pressure  causes  the  solution  to  boil  at  a  lower 
temperature  and  there  is  much  less  decomposition  of  the 
sugar  than  if  the  solution  were  boiled  down  at  atmospheric 
pressure.  The  crystals  of  sugar  which  are  deposited  on 
cooling  the  concentrated  solution  are  separated  from  the 
syrup  (molasses)  by  means  of  a  centrifugal  strainer. 

The  sugar  from  beets,  when  properly  purified,  is  identical 
with  cane  sugar. 

Maple  sugar  is  allowed  to  retain  substances  which  give 
to  it  a  desirable  flavor.  The  pure  sugar  is  the  same  as  that 
in  sugar  cane  or  sugar  beets. 


SUGARS  223 

Invert  Sugar. — If  ordinary  sugar  is  warmed  for  a  short 
time  with  a  dilute  acid  it  takes  up  water  and  is  converted 
into  a  mixture  of  equal  parts  of  two  other  sugars,  glucose 
and  fructose.  The  same  change  can  be  effected  by  certain 
organic  ferments  and  in  other  ways: 

C12H22On  +  H20  =  C6H1206  +  C6H1206 

Sucrose  Glucose  Fructose 

Invert  sugar 

Sucrose  rotates  the  plane  of  a  ray  of  polarized  light  to 
the  right.  Glucose  also  rotates  the  plane  of  the  ray  to  the 
right  but  fructose  rotates  it  to  the  left  to  a  greater  degree 
at  ordinary  temperatures  and  for  this  reason  the  mixture 
causes  a  left-handed  rotation  and  it  is  often  called  invert 
sugar.  Invert  sugar  is  found  in  honey,  in  some  fruit 
juices  and  in  syrups  made  from  sugar  cane  or  sorghum. 
Owing  chiefly  to  the  fructose  which  it  contains,  it  is  sweeter 
than  ordinary  sugar. 

Glucose. — The  formation  of  glucose  by  the  hydrolysis  of 
cane  sugar  has  just  been  mentioned.  It  may  also  be  pre- 
pared by  the  hydrolysis  of  starch  by  boiling  it  with  dilute 
sulfuric  or  hydrochloric  acid.  Large  quantities  of  the  sugar 
are  manufactured  in  this  way  and  sold  in  cheap  candies  and 
in  syrups,  especially  in  the  syrup  known  as  corn  syrup. 
The  acid  is,  of  course,  neutralized  or  removed. 

In  the  disease  called  diabetes,  sugar  or  starch  is'  con- 
verted in  part  into  glucose  and  eliminated  from  the  body 
in  that  form.  This  has  given  rise  to  an  impression  that 
glucose  is  not  a  safe  article  of  diet.  Ordinary  sugar  is 
converted  partly  into  glucose  during  digestion  and  there 
seems  to  be  no  scientific  ground  for  thinking  glucose  any 
more  harmful  than  sucrose. 

Maltose. — When  starch,  which  has  been  boiled  to  rupture 
the  granules,  is  mixed  with  malt  and  warmed  to  65°-70° 
it  is  largely  converted  into  a  sugar  called  maltose,  which  has 


224  CARBOHYDRATES,  ALCOHOLS,  ETC. 

the  same  composition  as  cane  sugar,  Ci2H22On,  but  differs 
from  it  in  some  of  its  properties: 

2(C6H10O5)n  +  nH20  =  nCi2H22On 

Starch  Maltose 

Malt  is  prepared  by  moistening  barley  and  allowing  it  to 
germinate.  As  the  barley  sprouts  an  enzyme  called  diastase 
is  formed.  This  is  soluble  in  water  and  a  small  quantity  of 
it  will  convert  a  large  amount  of  starch  into  sugar.  Dias- 
tase is  one  of  a  considerable  number  of  organic  substances 
called  enzymes,  which  act  as  catalytic  agents. 

Dextrin. — If  starch  is  moistened  with  very  dilute  nitric 
acid  and  heated  for  sometime  at  120°  it  is  converted  into  an 
easily  soluble  substance  called  dextrin.  This  is  used  in 
making  mucilage  and  for  the  backs  of  postage  stamps. 

Pectose.  Pectin.  Jelly. — Fruits  of  nearly  all  kinds, 
especially  when  not  fully  ripe  contain  an  insoluble  sub- 
stance called  pectose.  When  the  fruits  are  boiled  with 
water  the  pectose  is  decomposed  and  yields  a  soluble  sub- 
stance called  pectin.  Pectin  forms  a  jelly  with  sugar,  in  a 
slightly  acid  solution.  In  making  jelly,  from  one-half  to 
three-fourths  of  a  cupful  of  sugar  is  added  for  each  cupful 
of  fruit  juice.  The  fruit  juice  should  contain  from  0.5 
to  0.7  per  cent  of  acid,  calculated  as  tartaric  acid. "  The 
boiling  should  not  be  continued  too  long  after  separating 
the  juice  from  the  fruit,  as  the  pectin  seems  to  be  slowly 
destroyed  by  heat. 

Ethyl  Alcohol. — When  liquids  containing  sugar,  such  as 
the  juice  of  grapes,  apples  and  other  fruits,  or  syrups  ob- 
tained in  the  manufacture  of  sugars,  are  exposed  to  the 
air  they  almost  always  acquire  spores  of  yeast.  These 
grow  and  cause  the  fermentation  of  the  sugar  with  the  for- 
mation of  alcohol  and  carbon  dioxide.  The  fermentation 
seems  always  to  be  preceded  by  the  change  of  the  sugar  to 
invert  sugar  (see  cane  sugar,  above). 


ALCOHOL 


225 


The  solution  of  maltose  which  is  prepared  by  warming 
cooked  starch  with  malt  may  also  be  fermented  by  yeast  : 

C6H12O6  =  2C2H6O  +  2CO2 

Glucose  or  Alcohol 

fructose 


Ci2H22On 

Maltose 


H20  -  4C2H60  +  4C0 


In  manufacturing  alcohol,  corn  meal  or  potatoes  are  first 
thoroughly  cooked  and  the  cooked  material  is  mixed  with 


Dilute 
Alcohol  -»• 


Steam  Jacket- 


•*•  To  Condenser 


Alcohol- 
Free  Wafer 


FIG.  42. 


about  10  per  cent  of  its  weight  of  malt  and  enough  water 
so  that  the  resulting  solution  will  contain  about  10  per 
cent  of  sugar.  After  warming  for  a  short  time  to  bring 

15 


226  CARBOHYDRATES,  ALCOHOLS,  ETC, 

about  the  action  of  the  diastase  on  the  starch  the  solution 
is  cooled  to  about  the  temperature  of  the  hand  and  yeast  is 
added  and  the  mixture  is  allowed  to  ferment  for  three  or 
four  days.  The  fermented  liquid  is  then  distilled  to  sepa- 
rate the  alcohol  from  the  large  quantity  of  water  present. 
Alcohol  boils  at  78°  and  water  at  100°  and  when  a  mixture 
of  the  two  is  distilled  the  portions  passing  over  first  will 
contain  more  alcohol  than  the  original  liquid.  With  a 
" column"  still  constructed  on  the  principle  of  the  diagram, 
(Fig.  42),  by  introducing  the  dilute  alcohol  near  the  center 
the  stronger  and  stronger  alcohol  distils  upward  from  one 
shelf  to  another  while  the  water  containing  less  and  less 
alcohol  runs  downward  and  finally  leaves  the  still  practi- 
cally free  from  alcohol  at  the  bottom.  The  alcohol  which 
reaches  the  top  of  the  still  may  contain  90  per  cent  or 
more  of  alcohol  and  less  than  10  per  cent  of  water. 

Alcohol  is  used  for  burning,  as  a  solvent  in  making 
varnishes  and  in  preparing  pharmaceutical  extracts  and 
tinctures. 

"Denatured  alcohol"  contains  substances  which  have 
been  added  to  render  it  unsuitable  for  use  as  a  beverage.  It 
is  sold  free  of  tax  and  may  be  used  for  burning  and  in  making 
varnishes  but  it  is  much  more  poisonous  than  pure  alcohol 
and  is  unfit  for  drinking  or  for  any  medicinal  use. 

Acetic  Acid,  HC2H3O2. — When  dilute  alcohol,  such  as 
is  formed  by  the  fermentation  of  cider,  wine  or  other  sac- 
charine liquids,  is  exposed  to  the  air  in  loosely  closed  casks 
it  almost  invariably  acquires  bacteria  from  the  air,  which 
causes  a  second  fermentation  to  acetic  acid.  The  com- 
mercial product  is  called  vinegar.  Alcoholic  fermentation, 
which  is  caused  by  yeast,  takes  place  in  closed  vessels 
from  which  the  air -is  excluded.  The  acetic  fermentation 
requires  the  presence  of  air  to  furnish  the  oxygen  required 
for  the  oxidation  of  the  alcohol: 

C2H5OH  +  O2  =  C2H3O.OH(or  HC2H3O2)  +  H2O 


FATS,  SOAP  227 

In  vinegar  factories  beech-wood  shavings  are  inoculated 
with  the  bacteria  which  cause  -the  acetic  fermentation  and 
the  dilute  alcohol  is  permitted  to  run  slowly  over  these  in 
such  a  manner  as  to  expose  a  large  surface  of  the  liquid  to 
the  combined  action  of  air  and  the  bacteria.  Good  vinegar 
should  contain  4  per  cent  of  acetic  acid. 

Fats.  —  Such  substances  of  lard,  tallow,  olive  oil,  cotton 
seed  oil,  and  butter  are  composed  almost  entirely  of 
compounds  called  fats.  All  of  the  fats  contains  acids 
called  fatty  acids,  whose  hydrogen  has  been  replaced 
by  the  group  C3Hr,,  called  glyceryl.  This  group  is  charac- 
teristic of  glycerol  (commonly  called  glycerine)  C3H5- 
(OH)3.  The  three  most  common  compounds  found  in 
fats  are  derived  from  the  three  fatty,  acids,  stearic 
acid,  HCi8H35O2,  palmitic  acid,  HCi6H3iO2,  and  oleic 
acid,  HCisHssC^.  The  compounds  derived  from  these  are 
stearin,  C3H5(Ci8H35O2)3,  palmitin,  C3H5(Ci6H3iO2)3  and 
olein,  C3H5(Ci8H33O2)3. 

Soap.  —  When  a  fat  is  heated  with  a  concentrated  solution 
of  sodium  hydroxide,  NaOH,  it  is  decomposed,  forming 
the  sodium  salt  of  the  acid  and  glycerol,  C3H5(OH)3: 


2)3  +  3NaOH  =  3NaCi8H35O2  +  C3H5(OH)3 

Stearin  Sodium  stearate 

This  process  is  called  saponification  and  the  salts  formed 
are  called  soaps.  Soaps  dissolve  more  or  less  readily  in 
water  and  give  solutions  which  will  emulsify  fats  and  oils. 
if  olive  oil  or  kerosene  is  shaken  with  water  and  the  mixture 
is  allowed  to  stand  for  a  short  time  the  oil  separates  as 
a  layer  floating  on  the  water.  But  if  a  solution  of  soap 
is  added  and  the  mixture  shaken  again  an  emulsion  is  formed 
from  which  the  oil  and  water  will  separate  very  slowly  in- 
deed or  not  at  all.  On  rubbing  soiled  clothes  with  soap 
and  water  the  cleansing  of  the  cloth  depends  partly  on  the 
formation  of  an  emulsion  with  greasy  matters  on  the 


228  CARBOHYDRATES,  ALCOHOLS,  ETC. 

fabric,  because  the  soapy  solution  can  wet  such  substances 
and  emulsify  them,  while  pure  water  cannot.  The  emulsion 
and  particles  of  dirt,  which  are  also  wet  by  the  solution, 
can  then  be  rinsed  away  with  the  water. 

Glycerol  or  Glycerine,  C3H5(OH)3,  is  obtained  as  a 
by-product  in  making  soap  from  fats.  It  is  used  for  a 
variety  of  purposes  but  chiefly  in  the  manufacture  of  nitro- 
glycerine. 

Nitroglycerine  is  made  by  treating  glycerine  with  a 
mixture  of  nitric  and  sulfuric  acids. 

C3H5(OH)3  +  3HN03  =  C3H5(N03)3  +  3H2O 

Glycerine  Nitroglycerine 

Nitroglycerine  is  mixed  with  sawdust  or  some  other 
porous  substance  to  make  dynamite.  Either  nitroglycerine 
or  dynamite  is  exploded  by  a  detonating  cap  of  fulminate  of 
mercury.  The  explosion  is  due  to  the  combination  of  the 
oxygen  which  it  contains  with  its  carbon  and  hydrogen 
forming  a  large  volume  of  carbon  dioxide  and  steam  from  a 
small  volume  of  the  liquid.  The  combination  is  attended 
with  considerable  evolution  of  heat  and  this  causes  the 
expansion  of  the  gases  and  intensifies  the  explosion.  Some 
nitric  oxide,  NO,  is  formed  by  the  explosion  and  confined 
spaces  where  nitroglycerine  or  dynamite  is  exploded  require 
ventilation  before  entering  them  after  the  explosion. 

Phenol  or  Carbolic  Acid,  C6H5OH,  is  obtained  from 
coal  tar  and  is  also  prepared  synthetically.  It  is  used  as  a 
germicide  and  disinfectant  but  is  effective  only  when  applied 
directly.  The  vapor  is  not  concentrated  enough  to  be  of 
value.  Because  of  its  strong  disagreeable  odor,  uninformed 
persons  are  often  given  a  misleading  sense  of  security  by 
its  use.  It  is  the  chief  active  constituent  in  the  "  coal-tar 
dips"  used  in  the  care  of  sheep. 

Tartaric  Acid,  I&C^Oe—  When  grape  juice  is  allowed 
to  ferment  in  the  manufacture  of  wine  an  acid  salt  of  tar- 


BREAD  229 

taric  acid  which  is  known  as  cream  of  tartar,  KHC4H406, 
separates  because  the  salt  is  much  less  soluble  in  dilute 
alcohol  than  it  is  in  water.  The  chemical  name  of  the  salt 
is  acid  potassium  tartrate.  It  has  a  sour  taste  and  reacts 
readily  with  bases  or  with  carbonates,  giving  neutral  salts: 

KHC4H4O6  +  NaOH  =  KNaC4H4O6  +  H2O 

Sodium  potassium  tartrate 

KHC4H4O6  +  NaHCO3  =  KNaC4H4O6  +  CO2  +  H2O 

Bread. — The  manufacture  of  a  palatable  food  from  wheat 
flour  or  from  flour  prepared  from  other  cereals  depends 
largely  upon  securing  a  fine  cellular  structure  of  the  cooked 
material.  Such  a  structure  permits  the  easy  access  of  the 
saliva  and  other  digestive  fluids  to  all  parts  of  the  substance . 
and  in  this  way  promotes  its  digestion.  In  making  bread 
the  flour  is  mixed  with  water  or  milk,  or  both,  and  with  some 
yeast,  to  a  stiff  dough  and  the  whole  is  thoroughly  kneaded 
by  hand  or  by  some  mechanical  device  to  secure  the  fine 
cellular  structure  which  is  required.  The  yeast  acts  on 
the  small  amount  of  sugar  present  fermenting  it  to  alcohol 
and  carbon  dioxide.  The  latter  distends  the  little  cells  or 
interstices  in  the  dough,  causing  it  to  "rise."  After  a  time 
the  kneading  is  repeated  to  secure  a  more  uniform  structure 
and  after  being  allowed  to  rise  once  more  the  dough  is 
baked  in  an  oven.  It  is  important  that  the  materials  used 
shall  be  slightly  warm  and  that  the  bread  shall  be  kept  at  a 
temperature  of  21°-26°  (70°-80°  F.)  during  the  fermenta- 
tion or  rising.  At  a  lower  temperature  the  yeast  acts  too 
slowly  and  at  higher  temperatures  it  may  be  killed. 

The  alcohol  formed  by  the  fermentation  partly  escapes 
during  the  baking  of  the  bread  but  the  larger  portion  of 
it  is  retained.  The  amount  formed  is,  of  course,  small. 

Baking  Powders. — If  cream  of  tartar  (acid  potassium 
tartrate,  KHC4H4O6)  and  baking  soda  (acid  sodium  car- 
bonate, NaHCO3)  are  mixed  dry  no  reaction  occurs,  but 


230  CARBOHYDRATES,  ALCOHOLS,  ETC. 

on  adding  water  they  react  to  form  potassium  sodium 
tartrate,  carbon  dioxide  and  water  (see  tartaric  acid,  above). 
In  a  similar  manner  alum  (potassium  aluminium  sulfate, 
KM (80)4) 2),  will  not  react  with  baking  soda  when  dry,  but  on 
the  addition  of  water  carbon  dioxide  is  liberated  in  accord- 
ance with  the  equation: 


KAl(S04)->  +  3NaHCO3  =  KNaSO4  +  Na2S04  +3CO2  - 

Al(OH) 


Proteins.  Albumin.  Casein.  Gluten. — Complex  com- 
pounds containing  carbon,  hydrogen,  oxygen,  nitrogen  and 
usually  sulfur  or  phosphorus,  called  proteins,  are  found 
in  all  living  plants  and  animals.  The  most  familiar  of 
these  are  albumin,  which  forms  the  larger  part  of  the  white 
of  an  egg,  and  casein,  the  chief  constituent  of  the  curd  which 
can  be  separated  from  skim  milk.  Compounds  of  similar 
composition  are  found  in  the  muscular  fiber  of  meat  and  in 
the  gluten  of  wheat  flour,  which  remains  when  the  flour  is 
kneaded  between  the  fingers  in  a  stream  of  running  water, 
to  wash  away  the  starch. 

Digestion.  Formation  of  Tissues.  Production  of  Heat 
and  Energy. — Food  which  is  eaten  performs  in  the  body  two 
or  three  distinct  functions.  A  part  of  the  food  is  oxidized 
to  carbon  dioxide  and  water,  a  process  which  corresponds 
closely  to  the  burning  of  a  fire,  and  the  heat  generated 
maintains  the  temperature  of  the  body  above  the  tempera- 
ture of  the  air  which  surrounds  it.  A  part  of  the  energy  of 
the  food  is  also  converted  into  the  muscular  energy  with 
which  we  move  and  do  work.  If  these  were  the  only  func- 
tions to  be  served  in  the  body  our  food  might  consist 
exclusively  of  compounds  of  carbon,  hydrogen  and  oxygen, 
that  is,  of  such  compounds  as  the  fats  and  carbohydrates. 

1  Notice  that  enough  sodium  is  required  to  replace  the  aluminium, 
which  is  trivalent,  and  that  the  aluminium  forms  the  hydroxide,  Al(OH)s. 
The  rest  of  the  equation  follows  from  these  facts. 


GROWTH  OF  PLANTS  231 

It  will  be  seen  from  this,  too,  why  foods  containing  fats  are 
more  suitable  in  cold  than  in  warm  climates. 

The  tissues  of  the  body  are  also  constantly  broken  down 
and  must  be  restored  and  during  growth  new  tissues  must' 
be  produced.  Since  the  proteins  form  the  most  important 
part  of  the  tissues,  it  is  evident  that  foods  must  always 
contain  nitrogen,  sulfur  and  phosphorus  as  well  as  carbon, 
hydrogen  and  oxygen.  All  of  these  elements  must  also  be 
present  in  the  food  in  such  a  form  that  they  can  be  digested 
and  assimilated. 

In  the  process  of  digestion  proteins  are  decomposed  into 
simpler  compounds  which  are  soluble  and  which  can  then 
be  transported  by  the  blood  to  parts  of  the  body  where  they 
are  needed  for  building  or  restoring  tissues.  Fats  are 
emulsified  and  brought  into  the  circulation  in  that  form. 
By  respiration  the  oxygen  of  the  air  is  brought  to  one  side 
of  the  thin  membranes  of  the  lungs  while  the  venous  blood 
is  brought  to  the  other  side  of  the  membranes  by  circulation 
from  the  heart.  The  blood  brings  with  it  carbon  dioxide 
formed  by  the  oxidation  of  tissues  and  compounds  in  the 
body.  Through  the  membranes  of  the  lungs  it  gives  up 
this  carbon  dioxide  and  absorbs  oxygen  in  its  place.  The 
oxygen  is  carried  by  the  blood  through  the  arterial  circula- 
tion all  over  the  body  and  is  used  in  the  oxidation  which 
produces  heat  and  muscular  energy. 

Growth  of  Plants. — From  what  has  been  given  in  the  last 
paragraph  it  can  be  seen  that  the  animal  body  secures 
the  energy  to  maintain  its  existence  by  the  oxidation  of 
food.  The  energy  for  the  growth  of  plants  is  obtained  by  a 
radically  different  process.  The  carbon  for  the  plant  is 
taken  directly  from  the  carbon  dioxide  of  the  air.  The 
energy  to  separate  the  oxygen  from  the  carbon  of  the  carbon 
dioxide  is  absorbed  from  the  sunlight  by  the  leaves  of  the 
plant.  The  nitrogen  for  the  plant  must  be  furnished  by  the 
soil,  either  directly  in  the  form  of  nitrates,  or  ammonia,  or 


232  CARBOHYDRATES,  ALCOHOLS,  ETC. 

through  the  agency  of  nitrogen-fixing  bacteria  which  thrive 
in  the  roots  of  some  leguminous  plants,  such  as  clover  and 
alfalfa.  Potassium,  phosphorus,  sulfur  and  other  mineral 
constituents  necessary  for  the  growth  of  plants  must  also  be 
supplied  by  the  soil.  More  than  ninety  per  cent  of  the 
weight  of  growing  plants  is  derived  from  the  water  of  the 
soil  and  the  carbon  dioxide  of  the  air. 

Toxins  and  Antitoxins. — It  is  well  known  that  such  dis- 
eases as  diphtheria,  typhoid  fever,  tuberculosis,  yellow  fever, 
malarial  fever  and  many  others  are  caused  by  minute 
organisms  called  bacteria.  Some  of  these  organisms  when 
they  find  a  lodgment  and  grow  in  the  body  produce  poisons 
called  toxins.  These  are  often  very  virulent  and  may  pro- 
duce death.  It  has  been  discovered,  however,  that  under 
the  stimulus  of  the  toxin  the  body  produces  an  antidote 
called  an  antitoxin.  Very  little  is  known  about  the  exact 
nature  or  composition  either  of  the  toxins  or  antitoxins, 
but  it  has  been  discovered  that  in  some  cases  the  antitoxin 
may  be  developed  in  animals  and  used  as  an  antidote  for 
toxins  in  the  human  body.  Thus  by  inoculating  a  horse 
with  the  bacteria  which  cause  diphtheria  the  antitoxin  for 
the  disease  may  form  in  large  quantities  in  the  blood  of 
the  horse,  and  the  serum  from  the  blood,  if  injected  into  a 
person  suffering  from  the  disease  will,  in  most  cases,  effect 
a  cure. 

Alkaloids. — A  number  of  plants  produce  basic  substances 
containing  carbon,  hydrogen,  nitrogen  and  usually  oxygen, 
which  are  called  alkaloids.  This  name  is  given  to  them 
because  they  combine  with  acids  to  form  salts,  as  alkalies 
and  other  bases  do.  Some  of  the  alkaloids  are  powerful 
poisons  and  most  of  them  produce  marked  physiological 
effects.  The  best  known  are  strychnine,  morphine,  nico- 
tine, cocaine,  atropine  and  quinine.  Nearly  or  quite  all 
alkaloids  are  bitter;  some  of  them,  especially  strychnine 
and  quinine,  intensely  so. 


ALKALOIDS,  DYES  233 

Strychnine  is  a  very  powerful  poison  but  it  is  also  used 
in  small  doses  as  a  heart  stimulant. 

Morphine  is  obtained  from  the  poppy  and  is  used  to  pro- 
duce sleep.  It  is  the  chief  constituent  of  opium  and 
laudanum.  Morphine  and  opium  are  among  the  most 
dangerous  of  the  habit-forming  drugs. 

Nicotine  is  a  volatile,  liquid  alkaloid  found  in  tobacco. 
It  is  very  poisonous,  but  comparatively  small  quantities  of 
it  are  volatilized  in  smoking. 

Cocaine  is  used  to  produce  a  local  anesthesia  in  minor 
surgical  operations. 

Atropine  has  been  used  by  oculists  to  cause  a  widening 
of  the  pupil  of  the  eye  so  that  the  retina  may  be  examined  to 
better  advantage.  It  is  now  largely  displaced  by  other 
compounds  which  are  more  suitable. 

Quinine  is  obtained  from  Peruvian  bark.  It  is  specific 
for  malarial  fever  and  is  sometimes  used  for  other  pur- 
poses in  medicine. 

Dyes. — From  early  times,  natural  substances  obtained 
from  animal  and  vegetable  sources  have  been  used  to  give 
beautiful  colors  to  skins  and  cloths.  The  most  common 
natural  dyes  of  this  sort  are  cochineal,  Turkey  red,  indigo, 
fustic  and  logwood.  In  1856  a  young  English  chemist, 
William  H.  Per  kin,  discovered  that  a  beautiful  dye,  called 
mauve,  could  be  made  from  aniline,  a  compound  which 
can  be  prepared  from  the  benzene  of  coal  tar.  Not  long 
afterward  alizarin,  the  compound  which  gives  the  red  color 
of  Turkey  red,  was  prepared  artificially  from  another 
compound  found  in  coal  tar  and  it  was  very  soon  found 
that  the  artificial  alizarin  could  be  made  more  cheaply 
than  the  natural  product.  More  recently  indigo  has  been 
made  artificially  at  a  profit  and  since  1856  about  a  thousand 
dyes  which  are  made  artificially  have  been  patented.  Most 
of  these  are  different  from  any  of  the  natural  dyes  and 
almost  every  possible  shade  of  color  has  been  produced. 


234  CARBOHYDRATES,  ALCOHOLS,  ETC. 

Because  the  first  of  the  artificial  dyes  was  made  from 
aniline,  the  artificial  dyes  have  often  been  called  "  aniline 
dyes,"  a  name  which  is  not  very  correct.  The  designation 
"coal-tar  dyes"  is  more  proper,  as  nearly  all  of  the  artificial 
dyes  are  made  from  compounds  found  in  coal  tar.  Many 
of  the  dyes  at  first  discovered  fade  rapidly  on  exposure  to 
the  light  and  some  of  them  dissolve  or  "run"  when  the 
cloth  is  washed.  This  has  given  a  popular  impression  that 
the  artificial  dyes  are  inferior  to  those  from  natural  sources, 
but  this  is  by  no  means  true  of  all  of  them,  and  such  dyes 
as  alizarin  and  indigo  are  exactly  the  same  when  they  are 
made  artificially  as  when  obtained  from  madder  root  or 
the  indigo  plant. 

SUMMARY 

Carbohydrates  are  compounds  of  carbon  containing 
hydrogen  and  oxygen  in  the  same  proportions  as  in  water. 

Cellulose  is  the  principal  constituent  of  wood,  of  cotton, 
linen,  hemp  and  other  fibers,  and  of  paper. 

Nitrocellulose  is  used  in  gun  cotton,  smokeless  powder, 
lacquers,  collodion  and  celluloid. 

Starch  is  found  in  all  cereals  and  is  an  important  con- 
stituent of  foods  prepared  from  these. 

Sucrose  is  made  from  sugar  cane,  sugar  beets,  and  maple 
syrup.  It  is  changed  to  invert  sugar,  a  mixture  of  glucose 
and  fructose,  by  dilute  acids. 

Glucose  is  made  by  the  action  of  acids  on  starch.  It 
is  the  principal  constituent  of  corn  syrup. 

Maltose  is  made  by  the  action  of  the  diastase  of  malt  on 
starch. 

Dextrin  is  made  by  warming  starch  moistened  with  a 
little  dilute  nitric  acid. 

Pectose,  which  yields  pectin  on  boiling  with  water,  is 
found  in  most  fruits.  It  forms  a  jelly  with  acids  and  sugar. 


SUMMARY.     CARBOHYDRATES,  ETC.  235 

Ethyl  alcohol  is  made  by  the  action  of  yeast  on  liquids 
containing  maltose,  glucose  or  cane  sugar.  The  starch 
of  corn  or  potatoes  is  first  changed  to  maltose  by  the  use 
of  malt. 

Acetic  acid  and  vinegar  are  prepared  by  the  oxidation 
of  alcohol  under  the  influence  of  bacteria. 

Fats  contain  stearin,  palmitin,  olein  and  other  compounds 
of  organic  acids  with  glyceryl,  the  radical  of  glycerol. 
They  are  saponified  by  alkalies,  giving  soap  and  glycerol. 

Soap  helps  water  to  emulsify  greasy  substances  so 
that  they  can  be  removed. 

Nitroglycerine  is  glyceryl  nitrate.  Dynamite  is  a  mixture 
of  nitroglycerine  with  infusorial  earth  or  some  other  ab- 
sorbent material. 

Phenol  or  carbolic  acid  is  found  in  coal  tar  and  it  is  also 
manufactured.  It  is  used  as  a  germicide. 

Tartaric  acid  is  made  from  cream  of  tartar,  which  sepa- 
rates from  wines  during  fermentation. 

Bread  is  raised  by  yeast,  which  causes  the  fermentation 
of  sugar  to  alcohol  and  carbon  dioxide. 

Baking  powders  are  dry  mixtures  of  acid  sodium  car- 
bonate with  cream  of  tartar,  alum,  or  some  other  compound 
which  will  liberate  carbon  dioxide  from  the  baking  soda 
when  the  mixture  is  moistened. 

Albumin,  casein,  gluten  and  other  proteins  are  essential 
constituents  of  foods  and  are  required  for  restoring  tissues 
and  for  growth. 

In  digestion  foods  are  dissolved  and  partly  decomposed 
to  prepare  them  for  introduction  into  the  blood  and  for 
use  in  restoring  the  tissues  and  maintaining  the  warmth 
and  muscular  energy  of  the  body. 

Plants  utilize  the  carbon  dioxide  of  the  air  and  the  energy 
of  the  sun  and  store  energy.  Animals  dissipate  the  energy 
of  the  food  which  they  eat. 

Toxins  are  poisons  developed  in  the  progress  of  disease. 


236  CARBOHYDRATES,  ALCOHOL,  ETC. 

Antitoxins  are  antidotes  for  toxins  which  are  spontaneously 
developed  in  men  or  animals  affected  by  a  disease. 

Alkaloids  are  basic  compounds  found  in  plants.  The 
most  common  are  strychnine,  morphine,  nicotine,  cocaine, 
atropine  and  quinine. 

Dyes  were  formerly  mostly  of  vegetable  origin,  but  the 
larger  part  of  them  are  now  made  in  chemical  factories. 

EXERCISES 

1.  How  many  pounds  of  glucose  could  be  made  from  a  bushel 
of  corn  containing  60  per  cent  of  starch?     A  bushel  of  corn  weighs 
56  pounds 

2.  Assuming  that  90  per  cent  alcohol  has  a  specific  gravity  of 
0.79  and  that  a  gallon  of  water  weighs  8.3  pounds,  how  many 
gallons  of  90  per  cent  alcohol  could  be  made  from  a  bushel   of 
corn?     How  many  gallons  of  4  per  cent  acetic  acid? 

3.  Assuming  that  a  grease  used  for  soap  consists  of  palmitin, 
how  much  sodium  hydroxide  will  be  required  to  give  100  grams  of 
soap?  , 

4.  In  what  proportions  should  baking  soda  and  cream  of  tartar 
be  mixed  in  a  baking  powder?     If  tartaric  acid  were  used  in 
place  of  cream  of  tartar,  what  would  be  the  proportions  to  use? 

5.  What  weight  of  alum,  KA1(S04)2.12H20,  should  be  mixed 
with  100  grams  of  baking  soda  in  a  baking  powder? 


CHAPTER  XXI 
GROUP  IV:    SILICON,  TIN  AND  LEAD 

The  Carbon  Family  of  Elements. — Carbon  seems  to  be 
the  element  on  which  the  properties  of  the  living  matter  of 
plants  and  animals  chiefly  depend.  The  second  element  of 
the  same  family  is  silicon  and  this  is  equally  important  as 
the  most  abundant  element  after  oxygen,  in  the  solid  crust 
of  the  earth.  Silica,  the  dioxide  of  silicon,  SiO2,  and  sili- 
cates, salts  of  acids  of  which  silica  is  the  anhydride,  form 
a  very  large  proportion  of  the  soil  and  of  the  minerals  and 
rocks  which  are  most  abundant.  It  is  estimated  that  the 
element  silicon  forms  one-fourth  of  that  portion  of  the  earth 
which  we  have  been  able  to  examine.  Two  other  elements 
of  the  same  family,  tin  and  lead,  are  very  useful  common 
metals. 

Just  as  the  trioxides,  N203,  P203,  As2O3,  Sb203  and  Bi203, 
are  characteristic  of  the  elements  of  the  nitrogen  family, 
which  are  trivalent,  the  dioxides,  CO2,  SiO2,  PbO2  and 
SnO2  are  characteristic  of  the  quadrivalent  elements  of  the 
carbon  family. 

The  relations  among  the  atomic  weights  may  be  recalled 
by  the  following  table: 


C  12 

N  14 

O  16 

F  19 

Si  28 

P  31 

S  32 

Cl  35.5 

Ge  72 

As  75 

Se  78 

Br  80 

Sn  118 

Sb  120 

Te  127.6 

I  127 

Pb  207 

Bi  208 

— 

— 

While   arsenic,   antimony   and   bismuth   of  the  nitrogen 
family  are  brittle  and  in  some  of  their  properties  are  still 

237 


238  SILICON,  TIN  AND  LEAD 

closely  related  to  the  non-metals,  tin  and  lead  are  malleable 
and  fairly  good  conductors  of  heat  and  electricity.  Anti- 
mony and  bismuth  are  sometimes  called  half-metals,  but  tin 
and  lead  are  always  classed  as  metals. 

Occurrence  of  Silicon.  Preparation. — There  is  some 
reason  for  believing  that  all  of  the  carbon  on  the  earth  has 
been  at  one  time  combined  with  oxygen  and  that  we  find 
carbon  in  the  free  state  in  coal,  graphite  and  the  diamond 
because  of  the  power  which  plants  have  had  to  reduce  the 
carbon  dioxide  of  the  atmosphere  and  form  compounds 
which  have  since  decomposed  with  the  separation  of  free 
carbon.  Silicon  dioxide  is  not  reduced  in  the  growth  of 
plants  and  in  spite  of  the  abundance  of  the  element  there 
seems  to  be  no  process  occurring  in  nature  by  which  silicon 
is  liberated  in  the  free  state.  It  is  found  exclusively  as 
silicon  dioxide,  SiO2,  and  in  silicates  formed  by  the  union 
of  silicon  dioxide  with  other  elements. 

Silicon  is  prepared  commercially  by  heating  silicon  dioxide 
with  carbon  in  an  electric  furnace : 

SiO2  +  2C  =  Si  +  2CO 

Commercial  silicon  is  gray  and  crystalline.  It  is  used  in 
the  manufacture  of  ferrosilicon,  a  very  hard  form  of  iron, 
which  is  almost  insoluble  in  acids. 

Carborundum,  SiC,  is  also  prepared  by  heating  a  mixture 
of  silicon  dioxide  and  carbon  in  an  electric  furnace.  It 
is  a  crystalline  compound  very  much  harder  than  emery  and 
has  largely  replaced  the  latter  for  grinding  and  cutting 
purposes. 

Silicon  dioxide  is  found  abundantly  in  nature  in  a 
great  variety  of  forms.  The  clear,  transparent  forms  are 
called  rock  crystal.  Other  forms,  such  as  jasper,  amethyst, 
agate,  rose  quartz,  smoky  quartz  and  the  like  are  colored  by 
minute  amounts  of  other  substances.  The  mineralogical 
name  for  all  common  forms  is  quartz.  Opal  contains  a 


GLASS  239 

little  water.  Sand  usually  contains  a  large  proportion  of 
quartz,  because  it  is  harder  and  less  acted  upon  by  water 
than  the  other  minerals  in  the  rocks  from  which  the  sands 
have  been  formed. 

Quartz  may  be  fused  at  a  very  high  temperature  with 
the  oxyhydrogen  flame  or  in  an  electric  furnace  and  can 
be  made  into  tubes,  dishes  and  other  forms  of  apparatus 
which  are  useful  in  the  laboratory.  Such  apparatus  is 
much  less  soluble  than  glass  in  water  and  also  changes  so 
very  little  in  volume  when  heated  that  it  may  even  be 
heated  red  hot  and  quenched  in  water  without  cracking. 

Silicates. — Nearly  all  of  the  very  common  rocks  are 
composed  of  mixtures  of  minerals  which  are  silicates. 
Many  of  them  also  contain  quartz.  The  silicates  are 
derived  from  a  series  of  hypothetical  silicic  acids  none  of 
which  are  certainly  known  as  definite  compounds.  They 
all  have  the  same  anhydride,  silicon  dioxide.  Granite  is  a 
mixture  of  quartz,  mica  and  feldspar.  Other  common 
silicates  are  kaolin,  the  base  of  clay,  garnet,  talc  or  soapstone, 
asbestos  and  meerschaum. 

Glass. — When  silicon  dioxide  is  heated  with  the  oxides 
or  carbonates  of  such  metals  as  potassium,  sodium,  calcium, 
lead  and  some  others,  complex  silicates  are  formed  and  if 
the  proportions  are  properly  chosen  the  mixture  solidifies 
on  cooling  to  a  clear,  transparent  glass.  The  glass  made  in 
this  way  is  a  mixture  of  silicates  and  has  no  definite  melting 
point.  When  heated  it  softens  to  a  viscous  liquid  which 
can  be  easily  blown  into  large  bulbs  or  cast  or  molded  into 
various  forms.  As  it  does  not  crystallize  on  cooling,  but 
remains  clear,  it  is  suitable  for  the  great  variety  of  uses 
familiar  to  everyone.  Window  glass  and  the  glass  used  for 
bottles  and  other  common  articles  is  a  silicate  of  sodium  and 
calcium  with  small  quantities  of  other  metals  which  are 
present  as  impurities  in  the  sand,  lime  and  sodium  carbon- 
ate used  in  the  manufacture.  Flint  glass  contains  lead 


240  SILICON,  TIN  AND  LEAD 

in  place  of  calcium.  It  melts  at  a  lower  temperature,  is 
softer  and  has  a  higher  index  of  refraction.  Many  different 
kinds  of  glass  are  made  for  special  purposes  such  as  gage 
glasses  for  steam  boilers,  beakers  and  flasks  for  laboratory 
use,  thermometers,  lenses  and  imitations  of  diamonds  and 
other  precious  stones. 

Soluble  glass  is  a  silicate  of  sodium  prepared  by  fusing 
sand  and  sodium  carbonate: 

Na2CO3  +  SiO2  =  Na2SiO3  +  CO2 

Sodium  carbonate  Sodium  silicate 

All  kinds  of  glass  dissolve  in  water  to  a  slight  degree, 
though  the  ordinary  forms  are  practically  insoluble. 
Sodium  silicate,  however,  dissolves  to  a  large  extent, 
giving  a  viscous  solution  which  is  used  to  cement  glass 
and  porcelain,  to  fireproof  cotton  goods  and  to  preserve 
eggs.  It  is  decomposed  by  acids,  giving  silicic  acid  which 
may  remain  in  solution  as  a  colloid  or  may  be  precipitated 
in  a  gelatinous  form  according  to  the  way  in  which  the 
acid  is  added. 

Colloidal  Solutions. — Many  substances  which  are  usually 
insoluble  may  be  obtained  suspended  in  water  in  such  fine 
particles  that  they  do  not  settle  out  as  precipitates.-  This 
seems  to  be  partly  because  the  particles  are  so  very  fine 
(from  six  to  sixty  millionths  (0.000006  to  0.00006  mm.) 
of  a  millimeter  in  diameter)  and  partly  because  they  are 
each  composed  of  a  number  of  molecules  of  the  insoluble 
substance  gathered  about  some  positive  or  negative  ion. 
The  electrical  charges  of  the  ions  cause  the  particles  to 
repel  ;i«h  other  and  prevent  them  from  uniting  to  form 
large »  articles  which  would  fall  through  the  solution  as  a 
precipitate.  Solutions  of  this  character  are  called  colloidal 
solutions.  If  the  solution  of  colloidal  silicic  acid  is  placed 
in  the  parchment  sack  of  the  apparatus  shown  in  Fig.  43 
and  water  is  allowed  to  flow  slowly  through  the  bottle 
around  the  sack,  the  salt,  NaCl,  and  hydrochloric  acid, 


COLLOIDS.     DIALYSIS 


241 


FIG.  43. 


HOI,  of  the  solution  will  diffuse  through  the  parchment  and 

be  carried  away  while  the  silicic  acid  will  remain  behind. 

In  this  way  a  nearly  pure  solution  of  colloidal  silicic  acid 

may  be  obtained.     Such  a  process  is  called  dialysis  and  is 

often    used   to    separate    colloids   from   substances   which 

dissolve  in  the  molecular  or  ionic  condition.     The  latter  are 

sometimes  called  crystalloids.     The  distinction  seems  to  be 

due    to    the    fact    that    the 

colloids  consist  of  much  larger 

particles  and  for  that  reason 

diffuse  much  more  slowly  and 

may  also  be  unable  to  pass 

through  the   very   fine  pores 

of    the    parchment.     During 

the  process  of  digestion  food 

passes  partly  into  a  colloidal 

solution   and  the   colloids  of 

such  solutions  are  separated 

from  simple  molecalar  compounds,  which  are  also  formed, 

by  the  membranes  of  the  digestive  tract. 

The  very  fine  particles  of  clay  form  colloidal  solutions 
which  render  the  clay  plastic  and  make  it  possible  to  mold 
the  clay  into  the  forms  desired  for  the  manufacture  of 
brick,  tile,  earthenware  and  china.  Many  other  illustra- 
tions might  be  given  of  the  importance  of  colloidal  solutions 
for  natural  and  industrial  processes. 

Occurrence  of  Tin. — Tin  is  found  hi  nature  as  the  min- 
eral cassiterite,  the  dioxide,  SnO2.  Before  the  nineteenth 
century  the  world's  supply  came  almost  exclusively  from 
Cornwall,  England,  where  the  ores  have  been  mined  ever 
since  the  old  Roman  times.  Tin  is  now  obtained  from 
Banca,  the  East  Indian  islands  and  Tasmania.  Only 
small  quantities  have  been  found  in  the  United  States. 
Metallic  tin  is  readily  obtained  from  cassiterite  by  heating 
it  with  charcoal  or  coke, 

16 


242  SILICON,  TIN  AND  LEAD 

Tin  is  a  soft,  white  metal  which  melts  at  a  low  tem- 
perature. It  is  less  affected  than  any  other  common  metal 
by  water,  or  by  the  combined  action  of  water  and  air, 
and  for  this  reason  it  is  used  as  a  protective  coating  for 
sheet  iron  for  the  manufacture  of  tin  ware. 

Solution  Pressure  of  Metals. — If  a  piece  of  a  metal  is 
dipped  in  water  or  a  dilute  acid  or  alkali,  atoms  of  the  metal 
tend  to  pass  into  the  solution  as  positive  ions.  Thus 
iron  will  give  ferrous  ions,  Fe++,  tin  will  give  stannous  ions, 
Sn++,  zinc  will  give  zinc  ions,  Zn++.  As  these  positive 
ions  pass  into  the  solution,  the  piece  of  metal  becomes  nega- 
tive because  of  the  electrons  (atoms  of  negative  electricity) 
given  up  by  the  atoms  as  they  change  to  positive  ions, 
while  the  solution  becomes  positive.  The  amount  of  the 
difference  of  potential  between  different  metals  and  the 
same  solution  or  between  the  same  metal  and  different 
solutions  varies  with  the  character  of  the  metal  and  that 
of  the  solution.  Unless  there  is  some  way  provided  for 
the  electrons  to  escape  from  the  metal  the  escape  of  posi- 
tive ions  from  its  surface  is  almost  instantly  stopped  by 
this  difference  in  potential.  If  two  metals,  such  as  iron 
and  tin,  are  placed  in  water  or  a  dilute  acid,  the  difference 
in  potential  between  the  metal  and  the  solution  will  be 
greater  for  the  iron  than  for  the  tin;  in  other  words,  the 
iron  will  acquire  a  greater  negative  charge  than  the  tin. 
If  the  two  metals  are  connected  with  a  wire,  the  excess 
of  electrons  in  the  piece  of  iron  will  pass  into  the  tin. 
New  atoms  of  iron  may  then  pass  into  solution  till  the 
former  difference  of  potential  is  restored.  If  there  are 
tin  ions  in  the  solution,  electrons  will  escape  from  the  tin 
to  these  and  discharge  them,  leaving  the  tin  as  metallic  tin 
on  the  surface.  Under  these  conditions  the  stream  of 
electrons  will  be  continued  through  the  wire  from  the  iron 
to  the  tin  and  positive  ions  will  make  their  way  through 
the  solution  toward  the  tin.  In  this  manner  an  electric 


ELECTROMOTIVE  SERIES  243 

current  can  be  maintained  through  the  wire,  the  force 
which  drives  the  current  coming  from  the  fact  that  the 
solution  pressure  of  the  iron,  that  is,  the  tendency  of  the 
iron  to  pass  into  solution,  is  greater  than  the  solution 
pressure  of  the  tin. 

The  characteristics  of  tin  and  iron  which  have  just  been 
described  have  an  important  application  in  explaining 
the  conduct  of  the  tinware  used  in  the  household.  Pure 
tin  is  scarcely  affected  by  water  or  by  the  materials  usually 
employed  in  cooking.  Iron,  on  the  other  hand,  rusts  slowly 
in  contact  with  water  and  air.  A  piece  of  iron  in  contact 
with  a  piece  of  tin  will  rust  more  rapidly  than  the  iron 
alone  because  of  the  continual  escape  of  electrons  from  the 
iron  to  the  tin.  A  new  piece  of  tinware  in  which  the  coating 
of  tin  over  the  iron  is  perfect  does  not  rust  and  is  not  affected 
by  the  ordinary  liquids  used  in  cooking.  The  moment  a 
little  of  the  surface  of  the  iron  is  exposed,  however,  the  iron 
rusts  more  rapidly  than  if  the  tin  were  not  there.  This 
is  in  very  marked  contrast  with  the  conduct  of  iron  covered 
with  zinc  (p.  283). 

Electromotive  Series. — It  has  been  pointed  out  that  when 
iron  and  tin  are  put  in  a  solution  and  connected  by  means 
of  a  wire  a  stream  of  electrons  flows  from  the  iron  to  the 
tin  while  atoms  of  irqn  pass  into  solution  as  ferrous  ions, 
Fe++.  Because  the  iron  has  a  stronger  tendency  than  tin 
to  give  up  positive  ions  it  is  said  to  be  more  electro- 
positive than  tin.  Zinc  is  still  more  electropositive  than 
iron,  and  copper  is  less  electropositive  than  tin.  In  the 
following  table  the  common  metals  are  arranged  in  the  order 
of  the  potentials  which  they  assume  toward  each  other. 
Such  a  list  is  called  the  electromotive  series: 

Electropositive  end: 
Potassium 
Sodium 
Calcium 
Magnesium 


244 


SILICON-,.  TIN,  AND/  LEAD 


Aluminium 
Manganese 
Zinc 
Iron 
Nickel 
Lead 
Tin 

Hydrogen. 
Bismuth 
Copper 
Antimony 
Mercury 
Silver 
Platinum 
Gold. 
Electronegative  end. 


Direction  ofCvrre 


m\ 

- 


Zn 
Electro 

Positive 


FIG.  44. 


The  position  of  a  metal  in  the  electromotive  series  is  fixed 
by  determining  the  difference  in  potential  between  the 
metal  and  a  normal  solution  of  the  ions  of  the  metal. 

The  relative  positions  of  two  metals  in  the  series  may  be 
determined  by  conflteGting  stripe  of  the  metals  with  a  wire 


ALLOYS  OF  TIN  245 

and  determining  the  direction  in  which  the  current  flows 
(Fig.  44)  through  the  wire  when  the  strips  of  metal  are 
dipped  in  dilute  acid.  The  stream  of  electrons  (negative 
atoms  of  electricity)  flows  from  the  electropositive  metal 
through  the  wire  to  the  electronegative  metal,  i.e.,  in  such 
an  arrangement  the  electropositive  metal  is  negative  and 
the  electronegative  metal  is  positive.  Thus  if  strips  of  iron 
and  tin  are  used  the  iron  is  electropositive  and  the  positive 
current  of  electricity  flows  toward  it.1 

Alloys  of  Tin. — Besides  its  use  in  the  manufacture  of  tin- 
ware tin  is  used  in  a  number  of  alloys.  Solder  is  an  alloy 
of  tin  and  lead  and  a  similar  alloy  is  used  for  the  tin-plate 
for  roofing  purposes  because  lead  is  much  cheaper  than  tin. 
Babbitt  metal  is  mostly  tin,  lead  and  antimony  and  is  used 
in  the  bearings  of  machinery  because  of  its  antifriction 
properties.  Many  varieties  of  bronze  contain  tin  and  cop- 
per as  the  most  important  ingredients.  Tin  is  used  in  the 
coils  in  which  steam  is  condensed  in  preparing  distilled 
water  because  it  is  so  little  acted  on  by  water. 

Stannous  Chloride,  SnCl2.2H2O. — Tin  dissolves  readily 
in  concentrated  hydrochloric  acid  with  the  evolution  of 
hydrogen  gas.  The  metal  is  bivalent  in  the  chloride  which 
is  formed  but  it  has  a  strong  tendency  to  take  up  oxygen 
or  more  chlorine  and  pass  over  into  a  stannic  compound. 
This  makes  stannous  chloride  a  good  reducing  agent  and  it 
is  often  used  for  that  purpose. 

Stannic  Chloride,  SnCl4. — Tin  combines  directly  with 
dry  chlorine  to  form  stannic  chloride,  in  which  the  tin  is 

1  Confusion  sometimes  arises  in  the  minds  of  students  and  others  be^. 
cause  in  the  conventional  system  used  by  all  physicists  the  current  of 
electricity  is  represented  as  flowing  from  the  positive  pole  of  a  battery  or 
dynamo  to  the  negative  while  the  stream  of  electrons  flows  in  the  opposite 
direction.  When  the  older  theories  of  electricity  were  developed  there 
was  no  means  known  by  which  physicists  could  determine  the  true  direc- 
tion of  the  current  of  electrons  and  the  existence,  even,  of  the  electron 
was  not  suspected.  The  guess  which  was  made  as  to  the  direction  of  the 
current  was  the  opposite  of  the  truth  and  it  seems  impossible  now  to 
correct  the  mistake. 


246  SILICON,  TIN  AND  LEAD 

quadrivalent.  The  compound  is  a  volatile  liquid  which 
boils  at  114°.  As  iron  does  not  combine  so  readily  with 
chlorine  at  low  temperatures  and  ferric  chloride,  FeCl3, 
is  much  less  volatile,  scrap  tin  and  old  tin  cans  are  treated 
with  chlorine  gas  to  recover  the  tin.  The  stannic  chloride 
may  be  easily  reduced  to  metallic  tin  or  may  be  used  directly 
as  a  mordant  or  in  other  ways.  Stannic  chloride  fumes 
in  the  air  because  it  is  hydrolyzed  by  water,  giving  hydro- 
chloric acid.  This  shows  the  close  relationship  of  tin 
with  silicon  and  other  non-metallic  elements. 

Metastannic  acid,  H2Sn5Oii.9H2O,  is  an  insoluble  com- 
pound formed  when  tin  is  treated  with  nitric  acid.  Its 
formation  is  often  used  for  the  detection  and  determination 
of  tin  in  alloys,  since  all  of  the  other  metals  which  are 
acted  upon  by  nitric  acid,  except  antimony,  are  converted 
into  nitrates  by  the  acid  and  dissolve. 

Stannic  hydroxide,  or  stannic  acid,  Sn(OH)4,  is  formed 
by  precipitating  a  solution  of  stannic  chloride  with  ammo- 
nium hydroxide,  NH^OH.  It  dissolves  in  either  strong 
acids,  such  as  hydrochloric  or  sulfuric  acid,  or  in  strong 
bases,  such  as  sodium  hydroxide.  It  is,  therefore,  both  a 
base  and  an  acid.  The  solution  in  sodium  hydroxide  con- 
tains sodium  stannate,  Na2SnO3. 

Fireproofing  of  Cotton  Goods. — If  a  piece  of  cotton  cloth 
is  dipped  in  a  solution  of  sodium  stannate,  Na2SnO3, 
and  then,  after  squeezing  out  the  excess  of  the  solution  and 
drying  the  cloth,  it  is  dipped  in  a  solution  of  ammonium 
sulfate  and  again  squeezed  and  dried,  the  stannic  acid, 
H2SnO3,  or  stannic  oxide,  SnO2,  formed  by  these  processes 
combines  so  firmly  with  the  fiber  of  the  cloth  that  no  amount 
of  washing  will  remove  it.  Flannelette  and  other  forms  of 
cotton  cloth  which  have  been  treated  in  this  manner  will  no 
longer  catch  fire  and  burn.  If  the  process  could  be  generally 
introduced,  many  fatal  accidents  from  burning  might  be 
prevented.  (See  Professor  W.  H.  Perkin's  address  before 


LEAD  247 

the  International  Congress  of  Applied  Chemistry  in  1912.) 
In  some  states  laws  have  been  passed  requiring  the  fire- 
proofing  of  stage  curtains  in  schools  and  public  buildings. 

Lead,  Occurrence,  Metallurgy. — Lead  is  found  in  nature 
chiefly  in  the  form  of  galena,  PbS,  a  heavy,  black  mineral 
with  a  bright,  metallic  luster  on  fresh  surfaces.  The  mineral 
crystallizes  and  cleaves  in  cubes.  When  heated  in  the  air  it 
is  converted  into  a  mixture  of  sulfate,  PbSO4,  and  oxide,  PbO. 
When  this  mixture  is  heated  with  more  of  the  original  galena 
all  of  the  lead  is  reduced  to  the  metallic  state,  if  the  pro- 
portions are  properly  chosen: 

PbSO4  +  PbS  =  2Pb  +  2SO2 
2PbO  +  PbS  =  3Pb  +  S02 

Properties  and  Uses  of  Lead. — Lead  is  a  soft  metal  easily 
cut  with  a  knife  and  yielding  so  readily  to  pressure  that  it 
may  be  forced  through  a  die  of  the  proper  shape  into  the 
form  of  lead  pipe.  It  melts  easily  (327°)  but  at  a  consider- 
ably higher  temperature  than  tin  or  solder.  It  is  dissolved 
to  a  slight  degree  by  pure  water  and  is  not  suitable  for 
pipes  which  are  to  carry  water  for  household  use,  because 
lead  compounds  are  very  poisonous  and  even  a  very  minute 
amount  of  lead  taken  into  the  system  daily  through  a  series 
of  weeks  or  months  may  be  dangerous. 

Lead  is  very  slightly  attacked  by  dilute  sulfuric  acid,  even 
at  high  temperatures,  and  it  is  used  for  the  lead  chambers 
and  for  the  evaporating  pans  for  the  manufacture  of  sulfuric 
acid.  Lead  dissolves,  however,  in  hot,  concentrated  sulfuric 
acid.  The  conduct  of  iron  is  just  the  reverse.  It  dissolves 
easily  in  dilute  sulfuric  acid  but  is  scarcely  affected  by  the 
concentrated  acid.  Accordingly,  the  concentration  of  the 
dilute  acid  is  carried  on  in  leaden  pans  to  the  point  where 
the  lead  begins  to  dissolve  and  then  the  process  is  finished 
in  iron. 

Lead  is  used  for  bullets  and  shot  because  it  is  the  heaviest 


248  SILICON,  TIN  AND  LEAD 

of  the  cheap  metals  and  because  it  is  so  easily  given  the 
desired  forms.  Solder  is  an  alloy  of  lead  with  tin.  Babbitt 
and  antifriction  metals  are  alloys  with  tin  and  antimony, 
which  gives  hardness  and  sharpness  of  outline.  Lead  is 
also  a  constituent  of  the  fusible  alloys  used  for  safety  plugs 
in  steam  boilers,  in  electric  circuits  and  elsewhere. 

Oxides  of  Lead. — Lead  forms  three  important  oxides: 
litharge,  PbO,  red  lead,  Pb3O4,  and  lead  dioxide,  PbO2. 
Litharge  or  lead  monoxide  is  formed  when  lead,  or  galena, 
is  heated  to  a  rather  high  temperature  in  the  air.  Red 
lead,  Pb3O4,  is  formed  by  heating  litharge  or  white  lead  at  a 
somewhat  lower  temperature.  It  is  to  be  considered  as  a 
lead  salt  of  plumbic  acid,  H4PbO4,  and  the  formula  may  be 
written  Pb2Pb04,  showing  that  two-thirds  of  the  lead  is  in 
the  basic  condition  and  one-third  is  acid.  In  accordance 
with  this  formula,  on  treatment  with  nitric  acid  the  basic 
lead  dissolves  as  lead  nitrate,  Pb(NO3)2,  while  the  acid 
lead  remains  undissolved  as  lead  dioxide,  Pb02,  the  anhydride 
of  plumbic  acid,  H4Pb04.* 

Pb2Pb04  +  4HN03  =  2Pb(N03)2  +  Pb02  +  2H2O 

Red  lead  is  used  as  a  pigment.  It  will  be  noticed  that 
lead  dioxide  is  similar  in  formula  to  carbon  dioxide  and  that 
plumbic  acid  decomposes  into  lead  dioxide  and  water  as 
carbonic  acid,  H2COs,  decomposes  into  carbon  dioxide  and 
water. 

Storage  Batteries. — In  a  storage  battery  which  has  been 
charged  one  plate  contains  a  considerable  amount  of  spongy 
metallic  lead  and  the  other  plate  consists  largely  of  lead 
dioxide,  Pb02,  in  which  the  lead  is  quadrivalent.  The  jar  of 
the  battery  contains  dilute  sulfuric  acid.  It  will  be  remem- 
bered that  some  of  the  atoms  of  a  plate  of  metallic  lead  in 
contact  with  dilute  sulfuric  acid  lose  electrons  and  are 
changed  to  lead  ions,  Pb++: 

Pb  -  Pb++  +  2- 


STORAGE  BATTERIES  249 

These  lead  ions  combine  with  sulfate  ions,  SO4=,  of  the 
solution  to  give  insoluble  lead  sulfate,  PbSO4,  but  the  free 
electrons  remain  in  the  lead  plate  and  the  process  is  almost 
instantly  stopped  by  the  difference  in  potential  set  up 
between  the  plate  and  the  solution: 

Pb++  +  SO==PbS04 

If  the  lead  plate  is  connected  with  the  plate  containing 
lead  dioxide,  Pb(>2,  the  electrons  will  pass  to  that  plate 
and  combine  with  the  oxygen,  converting  it  into  oxygen  ions, 
0=,  and  leaving  the  lead  as  lead  ions,  Pb++. 

Pb02  +  2-  =  Pb++  +  20= 

The  lead  ions  will  combine  with  sulfate  ions  to  form  lead 
sulfate,  as  before,  while  the  oxygen  ions,  O=,  will  combine 
with  hydrogen  ions,  H+,  of  the  solution  to  form  water: 

0=  +  2H+  =  H2O 

At  the  end  of  the  discharge  both  plates  will  contain  lead 
sulfate.  In  charging  the  battery  an  external  electromotive 
force  is  applied  and  the  lead  sulfate  is  reduced  to  metallic 
lead  while  the  lead  sulfate  of  the  other  plate  is  oxidized  to 
lead  dioxide  with  the  liberation  of  sulfate  ions.  The  differ- 
ence in  potential  between  the  two  plates  is  about  two  volts 
and  it  is  evident  that  metallic  lead  and  lead  dioxide,  contain 
much  more  chemical  energy  than  the  equivalent  amounts 
of  lead  sulfate. 

The  sulfuric  acid  in  the  storage  battery  disappears  as  the 
battery  is  discharged  and  reappears  as  it  is  charged.  For 
this  reason  the  condition  of  the  cell  can  be  tested  by  deter- 
mining the  specific  gravity  of  the  liquid  in  the  cell.  Dilute 
sulfuric  acid  is  heavier  than  water  and  the  density  of  the 
liquid  will  approach  that  of  water  as  the  cell  is  discharged. 

Lead  nitrate,  Pb(NO3)2,  is  an  easily  soluble  salt  which 
can  be  prepared  by  dissolving  either  lead  or  lead  oxide 
in  nitric  acid. 


250  SILICON,  TIN  AND  LEAD 

Lead  chloride,  Pb.Cl2,  requires  125  parts  of  cold  water  for 
its  solution  and  is  formed  as  a  precipitate  when  hydro- 
chloric acid  or  a  chloride  is  added  to  a  solution  of  a  soluble 
lead  salt.  It  is  more  easily  soluble  in  hot  water. 

Lead  acetate  or  sugar  of  lead,  Pb(C2H3O2)2.3H2O,  is 
formed  when  litharge  is  dissolved  in  acetic  acid.  It  is 
easily  soluble  and  has  sometimes  been  used  for  hair  dyes, 
because  of  the  black  lead  sulfide,  PbS,  formed  when  it  is 
applied  to  the  hair.  Such  a  use  is,  however,  considered 
dangerous. 

White  lead,  one  of  the  most  valuable  white  pigments 
that  we  have,  is  a  basic  carbonate  of  lead  having  approxi- 
mately the  composition  2PbCO3.Pb(OH)2.  Thin  discs 
of  lead  are  placed  in  earthenware  pots  containing  acetic  acid. 
These  are  packed  in  series  with  layers  of  spent  tan  bark 
or  some  other  organic  material  which  will  ferment  and 
furnish  carbon  dioxide.  The  lead  slowly  corrodes  and  in 
three  or  four  months  it  is  almost  completely  converted  into 
solid,  brittle  cakes  of  the  basic  carbonate.  This  is  finely 
ground  and  mixed  with  linseed  oil  for  use  as  a  paint.  White 
lead  is  very  poisonous  and  great  care  is  needed  to  protect 
the  workmen  who  handle  it. 

Chrome  yellow,  PbCrO4,  is  a  brilliant  yellow  pigment 
used,  mixed  with  linseed  oil,  as  a  paint.  It  is  prepared 
by  precipitating  a  solution  of  some  soluble  lead  salt  with 
potassium  dichromate,  K2Cr207. 

SUMMARY 

The  carbon  family  of  elements  contains  carbon,  the  most 
important  element  of  living  bodies,  and  silicon,  the  most 
abundant  element  (except  oxygen)  in  minerals.  It  also 
contains  tin  and  lead. 

The  elements  of  the  group  all  form  dioxides,  and  all 
except  silicon  form  monoxides. 


SUMMARY.    SILICON,  TIN,  LEAD  251 

Silicon  is  found  chiefly  in  silica  and  in  silicates. 

Carborundum  is  a  very  hard,  crystalline  carbide  of  sili- 
con, used  as  an  abrasive. 

Silicon  dioxide  is  found  as  quartz,  jasper,  amethyst  and 
agate. 

A  great  variety  of  silicates  is  found  in  nature,  especially 
kaolin,  granite,  garnet,  talc  and  asbestos. 

Glass  is  an  artificial  silicate  of  sodium  or  potassium,  with 
calcium,  lead  or  other  metals.  It  is  viscous  through  a  wide 
range  of  temperature  and  remains  amorphous  and  trans- 
parent when  cold. 

Soluble  glass  is  a  silicate  of  sodium  which  will  dissolve 
in  water. 

A  colloidal  solution  contains  particles  much  larger  than 
ordinary  molecules  but  still  very  small.  These  particles 
remain  in  solution  either  because  they  are  too  small  to 
settle  out  or  because  they  cannot  coalesce  on  account  of 
electrically  charged  particles  which  they  contain. 

Colloidal  solutions  may  be  separated  from  electrolytes 
by  dialysis. 

Tin  is  found  as  cassiterite  in  Cornwall,  England,  the  East 
Indies  and  Tasmania. 

The  solution  pressure  of  a  metal  is  the  force  which  causes 
a  metal  to  give  positive  ions  to  a  solution  with  which  it  is 
in  contact  and  to  acquire,  for  that  reason,  a  negative  charge. 

The  electromotive  series  is  an  arrangement  of  metals  in 
the  order  of  the  magnitude  of  their  solution  pressures. 

The  most  important  alloys  of  tin  are  solder,  Babbitt  metal 
and  bronze. 

Tin  forms  two  chlorides,  stannous  chloride  and  stannic 
chloride. 

Stannic  hydroxide  is  both  a  base  and  an  acid.  It  is  used 
to  fireproof  cotton  goods. 

Metastannic  acid  is  formed  when  nitric  acid  acts  on  tin. 
It  is  insoluble. 


252  SILICON,  TIN  AND  LEAD 

Lead  is  found  as  galena  and  is  obtained  by  roasting  the 
mineral  and  heating  the  mixture  of  lead  sulfate,  lead  oxide 
and  lead  sulfide  which  is  formed. 

Lead  is  used  for  lead  pipe,  for  the  leaden  chambers  of  sul- 
furic  acid  works  and  in  solder,  type  metal,  Babbitt  metal 
and  safety  fuses. 

Lead  forms  three  oxides,  litharge,  red  lead  and  the  dioxide. 
Red  lead  is  lead  plumbate. 

Metallic  lead,  lead  dioxide,  and  sulfuric  acid  are  formed  in 
charging  a  storage  battery;  lead  sulfate  in  discharging  it. 

Lead  nitrate,  lead  chloride,  and  lead  acetate  are  common 
salts  of  lead. 

White  lead  is  a  basic  lead  carbonate  used  as  a  pigment. 

Lead  chromate  or  chrome  yellow  is  used  as  a  pigment. 

EXERCISES 

1.  How  much  silicon  dioxide  and  how  much  carbon  would  be 
required  to  prepare  a  ton  of  silicon? 

2.  How  much  of  each  will  be  required  to  prepare  a  ton  of  car- 
borundum? 

3.  If  a  colloidal  particle  is  spherical  and  has  a  diameter  of 
3 % oooooo  of  a  millimeter  and  a  specific  gravity  of  2,  how  many 
of  this  size  will  be  required  to  weigh  one  gram  ? 

4.  Design  a  die  which  might  be  used  to  press  lead  through  for 
the  manufacture  of  lead  pipe.     Draw  a  cross  section  of  the  top 
and  bottc  n  of  the  opening  through  which  the  lead  must  be  pressed. 
The  formation  of  the  tube  depends  on  the  welding  of  the  clean 
surfaces  of  lead  brought  together  with  the  die. 

5.  If  a  mixture  of  lime,  CaO,  and  litharge  is  heated  in  the 
air  calcium  plumbate  is  formed.     Write  the  equations  for  the 
reactions,    also    the  equations  for  the  reactions  which  will  occur 
when  calcium  plumbate  is  treated  with  dilute  hydrochloric  acid. 
In  what  proportions  should  the  litharge  and  lime  be  mixed? 

6.  Concentrated  hydrochloric  acid  converts  lead  dioxide  to 
lead  chloride;  what  other  products  will  be  formed?    Write  the 
equations.     What  other  dioxide  acts  in  a  similar  manner? 


CHAPTER  XXII 
GROUP  III:  BORON,  ALUMINIUM 

The  third  group  of  elements  contains  only  two  which  are 
used  either  in  the  metallic  state  or  in  their  compounds  for 
important  industrial  purposes.  These  are  boron  and  alu- 
minium. Boron  is  distinctly  non-metallic  in  its  properties 
and  is  practically  used  only  in  the  form  of  its  compounds. 
Aluminium  is  used  extensively  as  a  metal  and  several  of 
its  salts  are  also  important.  Each  element  is  trivalent. 

Boron,  Occurrence. — Boron  is  always  found  in  nature 
either  as  boric  acid,  H3B03,  or  in  the  form  of  salts  derived 
from  boric  acid  or  from  boric  anhydride,  B203.  Boric 
acid  issues  with  steam  from  fissures  in  the  ground  in  Tuscany, 
Italy,  and  considerable  quantities  of  the  acid  are  obtained 
from  this  source.  Borax  and  other  salts  of  the  boric  acids 
are  found  in  California,  Nevada  and  other  places  in  the 
west,  and  nearly  all  mineral  waters  contain  small  amounts 
of  borax. 

Borax,  Na4B2O7.10H2O. — Borax  may  be  considered  either 
a  salt  of  pyroboric  acid,  H4B207,  or  as  an  acid  salt  of  boric 
acid,  H3BOs.  In  the  latter  case  we  should  write  the  formula 
Na2Hio(BO3)4.  5H2O.  Boric  acid  is  so  very  weak  an  acid, 
however,  that  the  salt  shows  no  acid  properties  in  solution 
but,  on  the  contrary,  it  is  hydrolyzed  by  water  and  the 
solution  has  an  alkaline  reaction : 
Na  2Hi0(BO3)4  +  2H+  +  20H~  =  2Na+  +  4H3BO3  +  2OH- 

The  alkaline  reaction  is,  of  course,  due  to  the  hydroxide 
ions,-  OH~,  in  the  solution.  This  property  of  a  mild  alkali 
caused  by  its  hydrolysis  makes  it  suitable  for  washing 

253 


254  BORON,  ALUMINIUM 

flannels  and  other  delicate  fabrics  which  are  liable  to  injury 
if  sodium  hydroxide  or  some  other  strong  alkali  were  used. 

Borax  for  Welding.  Borax  Beads. — When  borax  is 
heated  it  swells  up  at  first  and  loses  water  but  finally  melts 
to  a  clear  glass  having  the  composition  Na2B4O7.  The 
formula  may  also  be  written  Na2(BO2)2.B2O3.  The  boric 
anhydride,  B203,  which  it  contains,  may  combine  with 
metallic  oxides  to  form  borates  exactly  as  silicon  dioxide, 
SiO2,  combines  with  lime,  litharge  or  other  oxides  in  the 
manufacture  of  glass.  Thus  if  borax  is  sprinkled  on  a  piece 
of  hot  iron  which  is  covered  with  a  coating  of  iron  oxide 
the  boric  anhydride  will  combine  with  the  oxide  to  form 
a  fusible  glass.  By  bringing  two  pieces  of  iron  together 
with  borax  between  them  the  film  of  oxide  on  the  surface 
of  each  is  dissolved  and,  on  pounding,  the  red-hot  surfaces  of 
pure  iron  unite  to  form  a  perfect  weld,  the  liquid  borax  being- 
forced  out  from  between  them.  Hot  borax  glass  will  dis- 
solve the  oxides  of  many  other  metals  and  some  of  these 
give  characteristic  colors  which  may  be  used  to  identify 
compounds  of  the  metals  in  the  laboratory. 

Boric  Acid,  HsBO3.  Borax  and  boric  acid  kill  or  prevent 
the  growth  of  bacteria  and  this  property  has  led  to  their 
use  in  preserving  food  and  as  an  eye-wash.  The  use  as  a 
food  preservative  is  now  forbidden  or  strictly  regulated  by 
federal  and  state  laws. 

Boric  acid  is  only  slightly  soluble  in  water.  It  is  prepared 
by  adding  hydrochloric  or  sulfuric  acid  to  a  warm  solution 
of  borax.  On  cooling,  the  boric  acid  crystallizes  from  the 
solution. 

Aluminium,  Occurrence. — Aluminium  is  the  third  ele- 
ment in  abundance  in  the  crust  of  the  earth,  oxygen  and 
silicon  being  first  and  second,  and  iron  fourth.  It  is  chiefly 
found  in  the  form  of  silicates.  A  large  proportion  of  the 
rocks  which  formed  the  crust  of  the  earth  in  the  earliest 
geological  time  must  have  consisted  of  silicates  of  alumin- 


ALUMINIUM  255 

him,  iron,  calcium,  sodium,  potassium  and  other  elements. 
During  many  millions  of  years  these  silicates  have  been 
broken  down  and  worked  over  by  the  action  of  water,  ice 
and  air,  with  the  aid,  in  some  cases,  of  vegetation  and  earth- 
worms. During  the  process,  partly  by  mechanical  agencies, 
partly  by  the  solvent  action  of  the  water,  the  original 
minerals  have  been  ground  to  exceedingly  fine  particles 
which  have  been  carried  away  by  the  water  and  deposited 
as  clay,  shales  and  sedimentary  rocks.  When  the  process 
is  most  complete  the  result  is  kaolin,  a  hydrated  silicate  of 
aluminium.  In  such  cases  the  sodium,  potassium  and 
other  elements  of  the  original  rocks  have  been  dissolved 
out  and  carried  away,  but  usually  the  clays  and  shales  are 
mixtures  of  kaolin  with  minute  particles  of  silica,  SiO2, 
and  of  other  minerals. 

Aluminium  is  also  found  in  the  oxide,  A1203,  as  the 
mineral  corundum.  In  its  crude  form  corundum  is  known 
as  emery  and  is  used  in  making  emery  wheels,  emery  paper, 
etc.,  because  it  is  the  hardest  natural  mineral  except  the 
diamond.  For  this  use  it  has  been  largely  replaced  by  car- 
borundum, SiC,  which  is  even  harder  than  corundum, 
though  less  hard  than  the  diamond.  Rubies  are  red  crystals 
of  corundum  colored  by  a  little  chromium.  They  are  now 
made  artificially.  Another  form  of  the  mineral  having  a 
blue  color,  is  sapphire. 

Metallic  aluminium  is  prepared  by  the  electrolysis  of 
aluminium  oxide,  A12O3,  dissolved  in  cryolite  or  some  other 
mineral  which  melts  at  a  low  temperature  and  contains  no 
water.  The  electrolysis  is  carried  out  at  such  a  temperature 
that  the  metallic  aluminium  collects  in  the  molten  form  in 
the  bottom  of  the  iron  pot  (Fig.  45)  which  is  used  for  the 
electrolysis.  It  is  drawn  out  from  time  to  time  through 
an  opening  in  the  end,  near  the  bottom. 

The  metal  melts  at  657°,  a  considerably  lower  tempera- 
ture than  the  melting  points  of  silver  and  copper ^  Its 


256 


BORON,  ALUMINIUM 


specific  gravity  is  2.6,  nearly  the  same  as  that  of  glass  and 
only  one-third  the  specific  gravity  of  iron.  The  metal  has 
a  bright,  silver- white  luster  and  does  not  tarnish  readily. 
For  this  reason  and  because  of  its  lightness  it  is  used  to  a 
considerable  extent  for  kitchen  utensils.  It  is  also  used 
for  telephone  wires,  and  in  an  increasing  number  of  valuable 
alloys  with  copper  and  other  metals.  Aluminium  dissolves 

readily  in  a  solution 
of  sodium  hydroxide, 
giving  sodium  alumi- 
nate,  Na3AlO3.  It 
also  dissolves  in  sul- 
furic  acid,  giving 
aluminium  sulfate, 
A12(S04)3. 

Goldschmidt's 
Thermite  Process. — 
When  aluminium  is 
burned  to  the  oxide, 

YIG.  45.  A12O3,  a  large  amount 

of    heat  is  liberated, 

much  more  than  by  the  burning  of  iron  or  other  common 
metals.  If  metallic  aluminium  is  mixed  with  ferric  oxide, 
Fe203,  and  the  mixture  is  ignited,  the  reaction  represented 
by  the  equation: 

2A1  +  Fe203  =  A12O3  +  2Fe 

takes  place  very  rapidly  with  the  evolution  of  such  a  quan- 
tity of  heat  that  the  whole  mass  is  heated  far  above  the 
melting  point  of  iron.  The  process  has  been  used  for 
welding  steel  rails,  repairing  broken  shafts  and  other 
similar  purposes. 

Aluminium  Sulfate,  A12(SO4)3.18H2O.— This  salt  is 
prepared  by  the  decomposition  of  clay  with  sulfuric  acid. 
It  is  easily  soluble  in  water  and  is  extensively  used  under 


'.I  I 


ALUM,  EARTHENWARE  257 

the  name  of  "alum"  for  the  clarification  and  purification  of 
water.  It  reacts  with  the  calcium  bicarbonate,  CaH2(CO3)  2, 
present  in  practically  all  natural  waters,  giving  a  precipi- 
tate of  insoluble  aluminium  hydroxide,  A1(OH)3,  which  is 
gelatinous  and  adheres  to  fine  particles  of  clay  and  to 
bacteria  in  the  water  in  such  a  manner  that  they  can  be 
removed  along  with  the  aluminium  hydroxide  bv^  filtration. 

AI2(S04)3  +  3CaH2(CO3)2  =  2A1(OH)3  +  3CaSO4  +  6C02 

Alum,  KA1(SO4)2.12H2O,  is  a  salt  with  an  astringent, 
sweetish  taste,  which  is  easily  prepared  by  crystallizing  a  mix- 
ture of  aluminium  sulfate  and  potassium  sulfate.  It  is  used 
extensively  in  the  cheaper  grades  of  baking  powders. 
For  this  purpose  it  is  mixed  with  "baking  soda, "  sodium 
bicarbonate,  NaHCO3.  The  alum  and  sodium  bicarbonate 
react  in  the  same  manner  as  the  aluminium  sulfate  and 
calcium  bicarbonate  in  the  purification  of  water. 

Alum  is  also  extensively  used  as  a  mordant  in  dyeing. 
The  aluminium  hydroxide  formed  by  its  hydrolysis  combines 
with  dyes  to  form  insoluble  compounds. 

Brick,  Earthenware,  Porcelain. — It  has  been  pointed  out 
that  ordinary  clays  are  mixtures  of  kaolin,  quartz  and  other 
minerals.  Some  of  these  minerals  melt  at  temperatures 
which  can  be  easily  obtained  in  ordinary  furnaces,  while 
quartz  and  pure  kaolin  require  very  high  temperatures  for 
their  fusion.  When  clay  has  been  molded  in  the  forms 
desired  for  brick,  tile,  earthenware  or  the  " biscuit"  of 
porcelain  it  may  be  "burnt"  in  such  furnaces,  the  "burn- 
ing" consisting  in  heating  the  mass  to  such  a  temperature 
that  a  part  of  the  minerals  present  melt  or  sinter  and 
bind  the  silica  and  other  minerals,  which  are  infusible  at 
the  temperature  used,  into  a  hard  solid  mass.  In  molding 
the  articles  the  clay  is  mixed  with  a  small  quantity  of  water 
which  forms  a  colloidal  solution  with  the  very  fine  particles 
of  the  clay  and  renders  the  mass  plastic. 

17 


258  BORON,  ALUMINIUM 

Materials  made  from  clay  in  the  manner  which  has  been 
described  remain  porous  and  are  not  suitable  for  domestic 
uses.  For  such  purposes  they  must  be  covered  with  a  glaze 
to  make  them  impervious  to  water.  The  glaze  is  sometimes 
formed  by  throwing  salt  into  the  furnace  toward  the  close 
of  the  burning.  The  salt  volatilizes  and  its  sodium  cojp- 
bines  with  the  silicon  and  other  elements  of  the  clay  to 
form  a  fusible  glass  while  the  chlorine  of  the  salt  combines 
with  the  hydrogen  of  the  water  present  in  the  gases  of 
the  furnace  and  escapes  as  hydrochloric  acid.  Porcelains 
are  usually  glazed  by  the  application  of  finely  ground  feld- 
spar. Glazes  containing  lead  are  often  used  for  the  cheaper 
kinds  of  earthenware. 

SUMMARY 

-  Boron  and  aluminium  are  the  only  common  elements  of 
Group  III. 

Boron  is  non-metallic,  aluminium  metallic.  Each  is 
trivalent. 

Boron  is  found  in  boric  acid,  borax  and  other  borates. 

Borax  is  a  salt  of  pyroboric  acid.  It  is  hydrolyzed  by 
water  and  is  used  for  laundry  purposes,  for  welding  of  iron 
and  in  the  detection  of  elements  by  the  colors  they  give  to 
borax  beads. 

Boric  acid  is  a  very  weak  acid  and  has  valuable  antiseptic 
properties. 

Aluminium  is  found  in  clay  and  in  many  silicates.  It  is 
obtained  by  the  electrolysis  of  a  solution  of  aluminium  oxide 
in  cryolite  or  some  similar  double  fluoride. 

Aluminium  is  a  light  metal  somewhat  resembling  silver 
in  appearance.  It  is  used  for  kitchen  utensils,  for  electric 
conductors  and  in  many  alloys. 

In  the  thermite  process  the  reaction  between  aluminium 
and  ferric  oxide  or  some  other  oxide  is  used  to  obtain  a  very 
high  temperature  for  welding  and  repairing  iron  or  steel. 


EXERCISES.     ALUMINIUM  259 

Alum  is  used  in  baking  powders  and  as  a  mordant. 

Aluminium  sulfate  is  used  as  an  aid  in  clarifying  and  puri- 
fying water. 

Brick,  earthenware  and  porcelain  are  made  by  the  partial 
melting  or  sintering  of  clay  or  kaolin. 

The  glaze  for  earthenware  and  porcelain  is  a  glass  con- 
taining feldspar  or  fusible  silicates  of  sodium,  lead  or  other 
metals. 

EXERCISES 

1.  In  what  proportions  should  aluminium  and  ferric  oxide  Lo 
mixed  for  use  in  the  thermite  process? 

2.  Many  alums  are  known  in  which  ammonium,  NH4,  sodium, 
Na,  or  some  other  univalent  metal  takes  the  place  of  potassium, 
and  others  in  which  iron,   Fe,   chromium,   Cr,   or  some    other 
trivalent  metal  replaces  the  aluminium.     Write  the  formulas  for 
the  following: 

Ammonium  alum  containing  aluminium  and  ammonium. 
Ferric  ammonium  alum  containing  iron  and  ammonium. 
Chrome  alum  containing  chromium  and  potassium. 
Rubidium  alum  containing  aluminium  and  rubidium. 

3.  The  best  known  baking  powders  contain  either  cream  of 
tartar  and  sodium  bicarbonate  mixed  with  starch,  or  alum  and 
sodium    bicarbonate    mixed    with    starch.     Calculate    the    per- 
centage composition  of  each  kind  of  powder,  assuming  40  per  cent 
of  starch  and  60  per  cent  of  the  other  ingredients.     How  many 
grams  of  carbon  dioxide  will  be  furnished  by  a  pound  of  each 
(1  Ib.  =  453  grams)?     How  many  liters  of  carbon  dioxide? 

4.  What  will  be  the  cost  per  pound  of  each  baking  powder  at 
the  following  wholesale   prices: 


Starch,  2^  cents  per  Ib. 

Alum,  2  cents  per  Ib. 

Cream  of  tartar,  26  cents  per  Ib. 

Sodium  bicarbonate,  3>£  cents  per  Ib. 


CHAPTER  XXIII 

GROUP  II,  FIRST  DIVISION:  ALKALI-EARTH  METALS, 
CALCIUM,  STRONTIUM,  BARIUM,  RADIUM 

Elements  of  the  Second  Group. — The  second  group  of 
elements  falls  into  two  very  distinct  divisions.  Magnesium 
and  calcium,  the  first  well-known  elements  of  the  group,  are 
rather  closely  related  and  strontium,  barium  and  radium 
resemble  calcium.  Zinc,  cadmium  and  mercury  of  the  sec- 
ond division  differ  very  markedly  from  the  metals  of  the 
first  division.  In  spite  of  some  resemblance  to  calcium, 
magnesium  is  usually  classified  with  the  second  division  of 
the  group. 

The  hydroxides  of  the  calcium  division  are  strong  bases 
and  the  metals  of  the  division  are  called  for  that  reason 
alkali-earth  metals.  The  hydroxides  of  zinc  and  cadmium  of 
the  second  division  are  less  basic  and  mercury,  the  last 
element  of  the  division,  forms  no  hydroxide. 

The  sulfates  of  the  alkali-earth  metals  are  less  and  less 
soluble  with  increasing  atomic  weights  and  the  extreme 
insolubility  of  radium  sulfate  is  used  in  separating  it  from 
other  elements. 

Calcium  ranks  fifth  in  abundance  among  the  elements  of 
the  earth's  crust  and  is  one  of  the  most  important  of  the 
elements.  Calcium  phosphate,  Ca3(PO4)2,  is  the  principal 
constituent  of  the  bony  skeleton  of  our  bodies  and  of 
the  bodies  of  all  vertebrate  animals.  Calcium  carbonate, 
CaCO3,  forms  the  skeleton  of  the  coral  insects  which  have 
built  the  immense  coral  reefs  in  the  ocean  and  of  insects 
which  have  built  the  limestones  which  were  formed  during 

260 


FLUORSPAR.     LIME  261 

millions  of  years  of  geologic  time.  Some  of  the  limestones 
have  been  metamorphosed  into  marble  by  heat  and  pres- 
sure. Some  of  them  contain  magnesium  carbonate,  MgCOs, 
as  well,  and  in  Switzerland  whole  mountains  are  made  up 
largely  of  dolomite,  MgC03.CaCO3,  a  mineral  containing 
calcium  carbonate  and  magnesium  carbonate  in  nearly 
equimolecular  proportions. 

Gypsum  is  a  soft,  crystalline,  hydrated  calcium  sulfate, 
CaSO4.2H2O,  which  is  called  alabaster  in  the  massive 
white  forms  used  for  vases  and  statuettes.  A  clear  crystal- 
line form  which  cleaves  in  thin  sheets  is  known  as  the  min- 
eral selenite.  Cruder  forms  are  used  in  the  manufacture  of 
plaster  of  Paris  and  as  a  fertilizer  to  furnish  sulfur  and  cal- 
cium to  soils  poor  in  these  elements  and  also  to  retain 
ammonia  in  the  soil. 

Fluorite  or  Fluorspar,  CaF2,  has  been  spoken  of  as  the 
chief  source  of  fluorine  and  its  compounds.  It  is  also  exten- 
sively used  as  a  flux  in  foundries  because  it  melts  at  a 
comparatively  low  temperature  and  dissolves  substances 
which  otherwise  might  remain  mixed  with  the  iron  and 
weaken  it. 

Calcium  is  now  easily  obtained  by  the  electrolysis  of 
melted,  anhydrous  calcium  chloride,  CaCl2.  It  is  a  white, 
crystalline  metal  which  decomposes  water  at  ordinary 
temperatures  with  the  formation  .of  calcium  hydroxide, 
Ca(OH)2.  It  is  sometimes  used  to  remove  traces  of  water 
from  absolute  alcohol. 

Calcium  Oxide  or  Lime  is  manufactured  on  a  large  scale 
by  heating  limestone,  CaCO3,  in  a  "  lime-kiln.7'  In  the  older 
'ime-kilns  the  pieces  of  limestone  and  coal  or  wood  were 
nixed  and  burned  in  some  form  of  furnace.  When  the 
materials  were  cool  the  lime  was  removed  from  the  furnace 
«vid  a  new  mixture  put  in.  At  the  present  time,  the  lime- 
stone and  fuel  are  charged  in  alternate  layers  into  the  top 
of  a  cylindrical  tower  and  the  finished  lime  is  raked  out  at 


262  ALKALI-EARTH  METALS 

the  bottom  from  time  to  time  without  stopping  the  process. 
This  method  is,  of  course,  much  more  economical  of  fuel. 

Calcium  Hydroxide  or  Slaked  Lime. — Calcium  oxide 
combines  with  water  directly  with  the  evolution  of  a  con- 
siderable quantity  of  heat : 

OH 
CaO  +  H2O  =  Ca(OH)2  or  Ca/ 

OH 

As  the  lime  combines  with  the  water  it  falls  to  a  very  fine 
'powder  and  if  a  considerable  excess  of  water  is  used  a 
milky  suspension  of  the  slaked  lime  in  water,  known  as 
milk  of  lime,  is  obtained.  If  a  larger  amount  of  water  is 
used,  the  calcium  hydroxide  settles  out  leaving  a  clear 
solution  of  the  hydroxide,  called  lime  water.  As  calcium 
hydroxide  is  only  slightly  soluble  in  water  the  solution  is 
always  very  dilute.  It  is  used  in  the  laboratory  for  the 
detection  of  carbon  dioxide.  Why? 

Mortar. — When  milk  of  lime  is  mixed  with  sharp,  clean 
sand  a  plastic  mass  is  obtained  which  is  used  between  layers 
of  brick  or  stone  in  building  walls  and  also  for  plastering  the 
walls  of  rooms.  As  the  water  is  absorbed  by  the  bricks  or 
dries  out  in  the  air  a  mass  of  considerable  initial  strength  is 
produced,  but  the  strength  is  greatly  increased  by  the  slow 
conversion  of  the  calcium  hydroxide  to  crystalline  calcium 
carbonate,  CaC03,  by  the  carbon  dioxide  of  the  atmosphere: 

Ca(OH)2  +  CO2  =  CaC03  +  H2O 

The  water  liberated  keeps  the  air  in  freshly  plastend 
rooms  moist  for  some  time.  The  crystals  of  calcium  rj,r- 
bonate  adhere  firmly  to  the  particles  of  sand  and  binj  the 
whole  together  to  a  solid,  hard  mass. 

Cement  is  manufactured  by  heating  to  a  high  temperatuio 
a  finely  powdered  mixture  of  limestone,  nearly  free  from 
magnesium  carbonate,  with  a  clay  rich  in  silica.  The  tern- 


HARD  WATERS  263 

perature  required  is  very  much  higher  than  that  of  a  lime- 
kiln as  the  silica  and  alumina  must  be  brought  into  combi- 
nation as  calcium  silicate  and  aluminate  and  this  requires 
the  sintering  of  the  mass  to  a  " clinker."  The  "clinker" 
is  again  finely  ground  and  mixed  with  a  little  plaster  of 
Paris,  giving  a  finished  cement  of  the  following  composition : 

Loss  on  ignition 0-2    per  cent 

Silica,  Si02 ! 15-20  per  cent 

Alumina,  A1203 3-8    per  cent 

Ferric  oxide,  Fe203 3-6    per  cent 

Lime,  CaO 58-64  per  cent 

Magnesia,  MgO 0-4    per  cent 

Potash  and  soda,  K2O,  Na20 0-2    per  cent] 

Sulfur  trioxide,  80s 0-2    per  cent 

When  mixed  with  water  the  cement  combines  with  it,, 
forming  partly  crystals  of  calcium  hydroxide,  Ca(OH)2, 
partly  hydrates  of  the  calcium  and  aluminium  silicates, 
which  "set"  to  a  hard  mass.  Sand,  gravel  or  other 
materials  are  added  to  increase  the  volume.  As  the  harden- 
ing is  not  due  to  carbon  dioxide,  cement  will  set  under 
water  and  is  often  called  "hydraulic  cement." 

Temporary  and  Permanent  Hardness  of  Waters. — Salts 
of  calcium  or  magnesium  render  water  "hard,"  that  is,  a 
considerable  amount  of  soap  must  be  used  with  such  water 
before  it  will  have  the  soft  feeling  characteristic  of  a  soapy 
water,  because  the  calcium  and  magnesium  combine  with 
the  fatty  acids  of  the  soap,  giving  a  curdy  precipitate,  and 
the  soap  cannot  produce  the  natural  effect  on  the  water 
till  all  of  the  calcium  has  been  precipitated. 

Many  natural  waters  contain  calcium  carbonate,  CaCO3, 
held  in  solution  by  carbonic  acid,  H2CO3,  as  calcium  bi- 
carbonate, CaH2(CO3)2.  When  such  a  water  is  boiled  the 
carbonic  acid  of  the  bicarbonate  dissociates  into  carbon 
dioxide  and  water  and  the  carbon  dioxide  escapes.  The- 


264  ALKALI-EARTH  METALS 

calcium  carbonate  which  remains  is  practically  insoluble 
in  water  and  separates  as  a  precipitate.  When  the  hard- 
ness of  the  water  is  due  to  calcium  bicarbonate  it  is  possible 
to  remove  it  by  boiling  the  water  and  allowing  the  calcium 
carbonate  to  settle,  and  this  kind  of  hardness  is  called 
temporary  hardness. 

Temporary  hardness  may  also  be  removed  by  adding  just 
the  right  amount  of  milk  of  lime,  Ca(OH)2.  Why? 

Calcium  sulfate,  CaSO4,  is  also  appreciably  soluble  in 
water  but,  as  it  is  not  held  in  solution  by  carbonic  acid, 
boiling  the  water  for  a  short  time  will  not  cause  the  calcium 
sulfate  to  precipitate.  For  this  reason  hardness  due  to 
calcium  sulfate  is  called  "  permanent  hardness. "  When 
the  water  is  boiled  away,  as  is  done  in  a  steam  boiler,  the 
calcium  sulfate  separates  and  produces  a  scale  which  adheres 
to  the  surface  within  the  boiler  and  is  particularly  trouble- 
some. Permanent  hardness  may  be  removed  by  the  ad- 
dition of  the  proper  amount  of  sodium  carbonate.  Why? 

Calcium  Sulfate,  Gypsum,  CaSO4.2H2O,  Plaster  of 
Paris,  2CaSO4.H2O. — When  gypsum  is  heated  it  loses 
part  of  its  water  and  is  converted  into  plaster  of  Paris,  which 
has  the  composition  2CaS04.H2O.  If  the  gypsum  is 
heated  to  too  high  a  temperature  all  of  the  water  may  be 
driven  out  but  the  anhydrous  calcium  sulfate,  CaS04, 
does  not  combine  readily  with  water  and  is  not  suitable 
for  the  uses  to  which  plaster  of  Paris  is  applied.  When 
plaster  of  Paris  is  mixed  with  water  to  the  consistency  of  a 
thick  cream  the  mixture  will  set  in  a  few  minutes  to  a  solid 
mass  of  gypsum,  CaSO4.2H2O,  as  the  plaster  combines 
with  the  water  and  the  mass  crystallizes.  The  material 
is  used  in  making  plaster  casts  of  works  of  art,  in  the  "hard 
finish"  for  plastered  walls,  and  for  many  other  purposes. 

Calcium  Chloride,  CaCl2. — A  solution  of  calcium  chloride 
is  easily  obtained  by  dissolving  marble  in  a  solution  of 
hydrochloric  acid.  When  the  solution  is  evaporated  the 


CALCIUM  PHOSPHATES  265 

hydrates  of  calcium  chloride  which  remain  retain  water  so 
firmly  that  it  is  necessary  to  heat  the  residue  to  250°  or 
above  before  all  of  the  water  is  expelled.  The  porous 
mass  which  remains  will  absorb  moisture  greedily  from 
ordinary  air  or  from  moist  gases  and  it  is  often  used  to  dry 
gases  in  the  laboratory.  If  exposed  to  ordinary  air,  the 
salt  will  finally  absorb  enough  water  to  deliquesce,  dis- 
solving in  the  water  which  it  attracts  to  itself. 

A  solution  of  calcium  chloride,  which  may  be  cooled  to 
20°-25°  below  0°  C.  without  freezing,  is  now  extensively 
used  in  refrigerating  machines  to  surround  the  cans  con- 
taining distilled  water  which  is  to  be  frozen  to  artificial  ice. 

OC1 

Chloride  of  Lime,  Ca<^        ,  has  been  described  in  con- 

Cl 
nection  with  hypochlorous  acid  (p.  86). 

Calcium  Phosphate,  Ca3(PO4)2. — The  occurrence  of  cal- 
cium phosphate  as  the  chief  mineral  constituent  of  bones 
has  been  already  mentioned.  Calcium  phosphate  is  also 
found  in  large  deposits  in  North  and  South  Carolina, 
Tennessee,  Georgia  and  Florida  and  is  extensively  used 
as  one  of  the  most  important  constituents  of  commercial 
fertilizers  which  are  applied  to  land  poor  in  available 
phosphorus.  As  calcium  phosphate  is  almost  insoluble 
in  water,  the  mineral  is  often  made  more  easily  soluble 
and  available  for  plants  by  treating  it  with  sulfuric  acid  to 
convert  it  into  the  acid  phosphate,  CaH4(PO4)2.  Tho 
mixture  of  calcium  sulfate  and  acid  calcium  phosphate  is 
called,  commercially,  a  "  super-phosphate. "  Write  the  equa- 
tion for  the  reaction.  Very  finely  ground  calcium  phosphate 
is  also  slowly  taken  up  by  plants  and  during  a  series  of 
years  may  be  as  useful  as  the  acid  phosphate. 

Calcium  Carbide,  CaC2,  is  manufactured  by  heating 
lime  and  coke  in  electrical  furnaces.  It  was  used  at  first 
for  the  preparation  of  acetylene  (p.  200)  for  illuminating 


266  ALKALI-EARTH  METALS 

purposes,  but  large  quantities  are  now  converted  into  cal- 
cium cyanamide  for  use  as  a  fertilizer. 

Calcium  Cyanamide,  or  "Lime -nitrogen,"  CaCN2. — At  a 
high  temperature  calcium  carbide  unites  with  nitrogen  to 
form  calcium  cyanamide  and  carbon: 

CaC2  +  N2  =  CaCN2  +  C 

Water  converts  calcium  cyanamide  to  calcium  carbonate 
and  ammonia: 

CaCN2  +  3H20  =  CaCO3  +  2NH3 

The  ammonia  formed  is  readily  available  in  the  soil  for 
the  growth  of  crops  and  for  this  reason  "lime-nitrogen" 
is  a  valuable  constituent  for  fertilizers  to  be  used  in  soils 
that  are  deficient  in  nitrogen. 

Atomic  Weight  of  Calcium.  Law  of  Dulong  and  Petit. — 
It  has  been  pointed  out  (p.  133)  that  the  most  satisfactory 
method  of  selecting  the  true  atomic  weight  of  an  element 
consists  in  finding  the  weight  of  the  element  contained 
in  a  gram-molecular  volume  (22.4  liters  at  0°  and  760  mm.) 
of  that  gaseous  compound  which  contains  the  smallest 
quantity  of  the  element  in  this  volume.  But  calcium  forms 
no  compound  whose  weight  in  the  gaseous  form  has  been 
determined,  and  a  considerable  number  of  other  elements 
form  no  compounds  which  can  be  converted  into  gases 
without  decomposition.  The  atomic  weights  of  such  ele- 
ments must,  of  course,  be  selected  in  a  different  manner. 
For  this  purpose  the  law  of  Dulong  and  Petit,  discovered 
in  1819,  has  been  useful.  These  chemists  found  that  the 
quantity  of  heat  required  to  raise  the  temperature  of  one 
gram  atom  of  an  element  one  degree  is  approximately  6.6 
calories.  If  this  quantity  of  heat  is  applied  to  7  grams  of 
lithium  or  to  65  grams  of  zinc  or  to  200  grams  of  mercury 
it  will,  in  each  case,  raise  the  temperature  one  degree. 

The  law  is  also  frequently  stated  that  the  specific  heat 


LAW  OF  DULONG  AND  PETIT 


267 


of  'an  element  multiplied  by  its  atomic  weight  is  a  constant 
quantity.     The  following  table  will  make  this  clear: 


Element 

Specific 
heat  . 

Atomic 
weight 

Atomic  heat 
(sp.  ht.  X  at.  wt.) 

Lithium 

0  94 

7  0 

6  6 

Graphite  (at  11°) 

0  16 

12  0 

1  9 

Graphite  (at  977°) 

0  467 

12  0 

5  6 

Silicon 

0  16 

28  4 

4  5 

Calcium 

0  17- 

40  0 

6  8 

Zinc 

0  093 

65  4 

6  1 

Bromine 

0  084 

80  0 

6  7 

Mercury 

0  033 

200  0 

6  7 

Lead 

0  03 

207  0 

6  2 

It  will  be  seen  from  the  table  that  graphite  and  silicon 
depart  rather  widely  from  the  law,  though  the  former 
approaches  it  more  closely  at  high  temperatures.  All  of 
the  metallic  elements  and  all  elements  having  atomic 
weights  above  40  conform  approximately  to  the  law.  The 
law  is,  at  best,  however,  only  approximate  and  is  of  service 
only  in  selecting  between  rather  widely  divergent  possible 
values  for  an  atomic  weight.  Thus  the  atomic  weight  of 
calcium  might  be  20,  40  or  60,  according  as  the  formula, 
of  the  chloride  is  CaCl,  CaCl2  or  CaCl3.  But  of  these  three 
values  only  an  atomic  weight  of  40  agrees  with  the  law. 
If  the  atomic  weight  were  20,  the  atomic  heat  (see  above) 
would  be  20  X  0.17  =  3.4.  If  it  were  60,  the  atomic  heat 
would  be  10.2.  An  atomic  weight  of  40  gives  an  atomic 
heat  of  40  X  0.17  =  6.8,  which  approximates  closely  to  the 
average  value  (6.6)  for  other  elements. 

The  laws  of  Avogadro  and  of  Dulong  and  Petit  have 
usually  been  considered  as  independent  and  wholly  un- 
related. l  A  little  consideration,  however,  shows  us  that  if 

1  See,  however,  G.  N.Lewis,  J.  Am.  Chem,  Soc.,  29, 1165  and  151G  (1907), 


268  ALKALI-EARTH  METALS 

we  accept  the  kinetic-molecular  theory  this  is  not  the 
case.  At  foundation  Avogadro's  law  depends  on  the  fact 
that  molecules  of  different  weights  exchange  energies,  when 
in  collision  with  each  other  or  with  the  walls  of  the  con- 
taining vessel  at  a  given  temperature,  in  such  a  manner 
that  the  average  value  of  ^mv2  (m,  mass,  v,  velocity)  is 
constant  and  is  independent  of  the  weight  of  the  molecule. 
The  law  of  Dulong  and  Petit  must  depend  on  a  similar 
property  of  the  atoms  of  the  elements  in  the  solid  or  liquid 
state. 

Strontium,  Sr,  and  Barium,  Ba,  are  found  as  sulfates, 
SrSO4  and  BaSC>4,  and  carbonates,  SrCO3  and  BaC03,  which 
are  all  nearly  insoluble  in  water.  It  will  be  noticed  that  the 
sulfate  and  the  carbonate  are  also  two  of  the  most  common 
compounds  of  calcium. 

Barium  Peroxide,  BaC>2,  is  formed  when  barium  oxide, 
BaO,  is  heated  in  the  air.  The  reaction  is  reversible  and 
at  a  higher  temperature  or  under  diminished  pressure  the 
peroxide  decomposes  into  barium  oxide  and  oxygen.  This 
property  has  been  much  used  in  the  preparation  of  oxygen 
gas.  Explain  how  and  write  the  equations. 

Barium  Chloride,  BaCl2.2H2O.— Barium  sulfate,  BaSO4, 
is  almost  the  only  salt  of  barium  which  does  not  dis- 
solve in  dilute  hydrochloric  acid.  For  this  reason  barium 
chloride  is  often  used  in  the  laboratory  for  the  detection 
of  sulfuric  acid  and  of  soluble  sulfates  and  for  the  determi- 
nation of  the  amount  of  these  which  may  be  present  in  a 
mixture  or  solution. 

Radium,  Discovery. — In  1896,  shortly  after  the  dis- 
covery of  Rontgen  or  X-rays,  Henri  Becquerel  in  Paris 
discovered  that  uranium  compounds  will  affect  a  photo- 
graphic plate  through  a  layer  of  black  paper  but  that  the 
effect  is  cut  off  by  metals  in  the  same  manner  as  the  X-rays. 
After  several  years  of  laborious,  painstaking  investigation 
Monsieur  and  Madame  Curie  showed  that  these  effects 


RADIOACTIVE  PROPERTIES 


269 


are  due  in  only  a  very  trifling  degree  to  pure  uranium 
and  that  they  are  caused  mainly  by  a  new  element,  radium, 
which  is  present  in  very  minute  quantities  in  minerals  which 
contain  compounds  of  uranium. 

Radium   and   its   compounds   exhibit   four   remarkable 
properties : 

1.  It  affects  photographic  plates  in  a  manner  similar  to 
the  effect  of  X-rays. 

2.  It  keeps  itself  at  a  higher  temperature  than  other  ob- 
jects around  it,  or,  in  other  words,  it 

is  constantly  giving  off  energy  in  the 
form  of  heat. 

3.  It  will  cause  air  in  its  neighbor- 
hood to  become  a  conductor  of  elec- 
tricity.      For    instance,    a    gold-leaf 
electroscope    (Fig.    4^),    which    will 
retain  its  charge  for  a  long  time  in 
ordinary  air  as  shown  by  the  repul- 
sion of  the  leaf,  will  be  quickly  dis- 
charged  and   the  leaf  will  fall  if  a 
mineral  containing  radium  is  brought 
near  to  it.     This  has  been  found  to 

be  the  most  accurate  method  of  detecting  and  measuring 
the  amount  of  radium  and  of  other  radioactive  elements 
in  minerals,  water,  air  or  other  substances. 

4o  Radium,  when  brought  near  to  the  skin  for  some  time,, 
may  produce  severe  burns,  somewhat  similar  to  sunburn, 
and  it  kills  bacteria  very  much  as  sunlight  does.  The  use 
of  radium  in  the  treatment  of  cancer  seems  closely  related 
to  this  property.  The  amount  of  radium  to  be  found  is  so 
small  and  its  value  for  medicinal  and  other  purposes  is  so 
great  that  compounds  of  the  element  have  been  sold  at  a 
price  of  over  $100,000  a  gram  of  radium,  while  gold  is  worth 
only  about  64  cents  a  gram. 

5.  Radium  causes  zinc  sulfide  and  some  other  substances 


FIG.  46. 


270  ALKALI-EARTH  METALS 

to  phosphoresce  and  glow  in  the  dark.  This  property 
of  radium  is  used  in  the  radiolite  dials  of  watches. 

Disintegration  of  Elements. — For  a  thousand  years  or 
more,  through  the  middle  ages,  a  class  of  men  called  alchem- 
ists sought  by  every  'means  they  could  think  of  to  find  some 
method  of  converting  lead  and  other  base  metals  into  gold. 
The  'failure  of  their  quest  for  the  " Philosopher's  Stone'7 
contributed  to  the  conclusion  finally  reached  by  chemists 
that  it  is  impossible  to  transform  one  chemical  element 
into  another.  This  conclusion  was  universally  accepted 
at  the  close  of  the  nineteenth  century.  In  spite  of  this 
belief  Professor  Rutherford,  who  was  then  working  at  Mc- 
Gill  University  in  Montreal,  proposed  the  startling  hy- 
pothesis that  radioactive  elements  disintegrate  with  the 
formation  of  other  elements.  Shortly  after  this  Soddy, 
who  began  with  Rutherford  and  who  continued  his  studies 
with  Professor  Ramsay  in  London,  demonstrated  that 
helium  is  formed  by  the  spontaneous  decomposition  of 
radium.  In  the  decomposition  the  helium  atoms  are  shot 
out  with  a  tremendous  velocity,  such  that  they  will  pene- 
trate thin  layers  of  ordinary  opaque  matter  and  affect  a 
photographic  plate  beneath.  They  also  produce  the  other 
radioactive  phenomena  referred  to  in  the  preceding 
paragraph. 

To  account  for  these  phenomena  it  is  supposed  that  the 
atoms  of  the  radioactive  element,  and  probably  of  all  other 
elements,  are  composed  of  positively  charged  particles 
around  which  electrons1  are  revolving  with  a  high  velocity. 
Occasionally  one  of  the  positive  particles,  having  the  weight 
of  a  helium  atom,  gets  into  an  unstable  position  with 
reference  to  the  other  parts  of  the  atom  and  is  shot  out. 
After  the  helium  atom  has  escaped,  the  residue  is  no  longer 
an  atom  .of  radium  but  it  is  now  an  atom  of  radium  ema- 
nation or  niton  (p.  158)! 

1  Atoms  of  negative  electricity. 


LIFE  OF  AN  ELEMENT  271 

It  is  evident  from  the  relation  of  radium  to  helium  and 
niton  that  radium  atoms  are  composite,  but  radium  cannot 
be  considered  as  merely  a  compound  of  helium  and  niton 
in  the  sense  in  which  we  speak  of  a  compound  of  sodium  and 
chlorine.  No  means  is  known  by  which  niton  and  helium 
can  be  reunited  to  form  radium  and  all  three  have  those 
properties  by  which  we  are  accustomed  to  distinguish 
elements  from  compounds. 

Having  discovered  that  radium  atoms  are  composite  it 
is  natural  to  suppose  that  the  atoms  of  the  other  elements 
are  composite,  also,  and  that  the  relations  of  the  elements 
in  the  periodic  system  are  due  to  this  fact.  An  extremely 
interesting  field  for  investigation  has  been  opened  up  in 
this  connection  and  there  seems  to  be  a  good  probability 
that  our  knowledge  of  the  structure  of  atoms  will  develop 
rapidly. 

Life  of  an  Element. — The  rate  at  which  the  atoms  of  dif- 
ferent radioactive  elements  disintegrate  varies  very  greatly. 
One-half  of  a  given  quantity  of  radium  would  decompose 
into  niton  and  helium  in  about  1800  years.  One-half  of 
the  remainder  would  decompose  in  the  next  1800  years, 
and  so  on.  It  is  evident  that  no  definite  time  can  be  given 
when  all  of  the  element  would  be  disintegrated,  but,  by 
common  consent,  the  time  required  for  the  disintegration  of 
one-half  of  its  atoms  is  spoken  of  as  the  "  half-life  period  of 
the  element."  Radium  is  formed  indirectly  from  uranium, 
but  while  the  half -life  period  of  radium  is  1800  years  that 
of  uranium  is  about  6,000,000,000  years.  If  radium  is 
formed  only  from  uranium,  as  seems  probable,  it  is  clear 
that  the  amount  of  radium  in  the  world  must  be  very  small 
in  comparison  with  the  amount  of  uranium.  The  half-life- 
period  of  niton  is  3.8  days  and  only  exceedingly  minute 
quantities  of  that  element  can  be  obtained.  A  cubic 
millimeter  of  niton  weighs  only  about  0.01  milligram,  but 
with  a  few  cubic  millimeters  of  the  gas  Professor  Ramsay 


272 


ALKALI-EARTH  METALS 


determined  its  boiling  point  and  density  and  showed  that 
its  atomic  weight  is  about  222.  The  atomic  weight  of 
radium  is  226.4  and  that  of  helium  is  4. 

Other  Radioactive  Elements. — In  the  disintegration  of 
radium,  as  has  been  stated,  helium  atoms  are  shot  out  and 
the  residue  which  remains  consists  of  niton,  an  element 
with  properties  widely  different  from  radium.  Niton,  in 
turn,  disintegrates  more  than  150,000  times  as  fast  as 
radium.  As  it  does  so  it  loses  another  helium  atom  and 
the  residue,  radium  A,  is  supposed  to  have  the  atomic  weight 
of  218.4,  but  the  element  is  known  only  by  its  radioactive 
properties.  Radium  A  also  decomposes  rapidly  but  shoots 
out  electrons  instead  of  helium  atoms.  As  an  electron 
weighs  only  one-eighteen-hundredth  part  as  much  as  a 
hydrogen  atom,  radium  B  has  practically  the  same  atomic 
weight  as  radium  C. 

Lead  is  supposed  to  be  the  end  of  the  series  of  products 
formed  from  radium.  The  relation  between  the  elements  of 
the  series  can  be  seen  from  the  following  table : 


Element 

Atomic 
weight 

Half-life  period 

Atoms  expelled 

Radium 

226  4 

1760  years 

Helium 

Niton 

222  4 

3  86  days 

Helium 

Radium  A  

218.4 

3  minutes 

Helium 

Radium  B  

214.4 

26  .  7  minutes 

Electrons 

Radium  C  

214.4 

19  .  5  minutes 

Helium  and  electrons 

Radium  D  
Radium  E  

210.5 
210.5 

17.3  years 
6  .  2  days 

None 
None 

Radium  E 

210  5 

4  8  days 

Electrons 

Radium  F  
Inactive     product, 

210.5 

ofifi    K: 

143  days 

Helium 

Besides  the  radium  series  of  radioactive  elements,  which 
begins  with  uranium  and  probably  ends  with  lead,  there  is 


COUNTING  MOLECULES  273 

a  second  series  beginning  with  thorium  and  supposed  to 
end  with  bismuth,  and  a  third  series  beginning  with  actin- 
ium, an  element  known  only  by  its  radioactive  properties, 
and  ending  in  an  inactive  product  which  cannot  be  identi- 
fied. The  atomic  weight  of  no  element  in  the  series  is  known. 

Chemists  of  the  Bureau  of  Mines  have  recently  proposed 
the  use  of  mesothorium  of  the  thorium  series  to  replace 
radium  for  the  dials  of  radiolite  watches. 

A  very  surprising  result  of  recent  investigations  is  the 
fact  that  the  density  and  atomic  weight  of  lead  from  radio- 
active sources  are  different  from  those  of  ordinary  lead,  al- 
though the  two  kinds  of  lead  are  alike  in  other  chemical 
and  physical  properties. 

Counting  the  Number  of  Molecules  in  One  Cubic 
Centimeter  of  a  Gas. — It  has  been  shown  that  one  gram  of 


FIG.  47. 

radium  gives  off  158  cubic  millimeters  of  helium  gas  in  a 
year.  It  would  take  a  little  more  than  six  years  for  it  to 
give  one  cubic  centimeter.  This  determination  made  it 
possible  for  Rutherford  and  Geiger  to  count  the  helium 
atoms  in  a  very  small  volume  of  the  gas  and  estimate  the 
number  in  one  cubic  centimeter.  As  the  helium  atom  and 
the  helium  molecule  are  identical,  on  the  basis  of  Avogadro's 
law  this  number  is  the  same  as  the  number  of  molecules  in 
a  cubic  centimeter  of  any  other  gas. 

The  apparatus  used  for  the  experiment  is  shown  in  Fig. 
47.     A  small  amount  of  radium  was  placed  in  A,  which  was 

18 


274  ALKALI-EARTH  METALS 

completely  exhausted.  A  thin  screen  of  mica  at  D  made 
it  possible  to  leave  a  little  air  in  C  while  the  helium  atoms 
shot  out  by  the  radium  would  pass  through  the  mica  into 
C.  E  was  a  charged  wire  connected  with  an  electrometer. 
When  a  helium  atom  shot  into  C  the  air  was  ionized  and 
became  a  conductor  for  electricity.  The  discharge  of  E 
through  the  air  to  the  walls  of  C  was  immediately  recorded 
by  the  electrometer.  The  amount  of  radium  and  its  distance 
from  the  opening  covered  by  the  mica  were  so  chosen  that, 
on  the  average,  ten  or  eleven  helium  atoms  passed  into  C 
in  one  minute.  The  radium  shot  out  the  same  number  of 
atoms  in  all  directions,  of  course,  and  by  measuring  the 
distance  of  the  radium  from  the  mica  screen  and  the  size 
of  the  opening  through  which  the  atoms  passed  into  C, 
it  was  easy  to  calculate  how  many  atoms  were  shot  out  in 
a  minute  or  an  hour  from  the  radium  used.  From  this 
number  and  from  the  time  required  for  the  amount  of 
radium  used  to  send  out  one  cubic  centimeter  of  helium  gas 
the  number  of  molecules  in  one  cubic  centimeter  was  calcu- 
lated. 

By  placing  a  phosphorescent  screen  of  zinc  sulfide  back 
of  the  mica  plate  a  flash  was  produced  by  each  helium  atom 
which  struck  the  sulfide.  By  counting  the  flashes  the 
number  of  helium  atoms  was  determined.  Both  methods 
gave  results  which  were  nearly  the  same  as  the  best  de- 
terminations of  the  number  of  molecules  in  a  cubic  centi- 
meter of  a  gas  which  have  been  made  by  other  methods. 
This  number  is  about  27,100,000,000,000,000,000. 

SUMMARY 

The  first  division  of  the  second  group  contains  calcium, 
strontium,  barium  and  radium — the  alkali-earth  metals. 

The  hydroxides  are  strong  bases,  the  carbonates  are 
insoluble  and  the  sulfates  are  difficultly  soluble  or  insoluble. 


SUMMARY.    ALKALI-EARTH  METALS  275 

Calcium  occurs  as  the  carbonate,  sulfate,  phosphate  or 
fluoride. 

Lime  is  manufactured  by  heating  calcium  carbonate. 

Calcium  hydroxide  is  used  in  mortar  and  plaster. 

Cement  is  a  mixture  or  compound  of  lime  with  calcium 
and  aluminium  silicates.  It  combines  slowly  with  water, 
forming  calcium  hydroxide  and  hydrates  which  set  to  a 
hard  mass. 

Calcium  sulfate  gives  "permanent"  hardness,  which 
cannot  be  removed  by  boiling  a  natural  water.  Calcium 
bicarbonate  gives  " temporary"  hardness,  which  can  be 
removed  by  boiling. 

Plaster  of  Paris  is  a  hydrate  of  calcium  sulfate  which  com- 
bines readily  with  more  water  to  form  another  hydrate  that 
sets  to  a  hard  mass. 

Anhydrous  calcium  chloride  is  used  as  a  drying  agent.  A 
solution  of  calcium  chloride  with  a  low  freezing  point  is 
used  in  refrigerating  machines. 

Chloride  of  lime  is  partly  calcium  chloride  and  partly 
calcium  hypochlorite. 

Calcium  phosphate  and  "super-phosphate"  are  used  as 
fertilizers. 

Calcium  carbide  is  used  for  the  preparation  of  acetylene, 
also  for  the  manufacture  of  calcium  cyanamide,  or  lime- 
nitrogen.  The  latter  is  used  as  a  fertilizer. 

Barium  peroxide  is  used  in  the  manufacture  of  oxygen 
and  of  hydrogen  peroxide. 

Barium  chloride  is  used  for  the  detection  and  determina- 
tion of  sulfuric  acid. 

Radium  is  formed  very  slowly  by  the  disintegration  of 
uranium.  It  disintegrates,  itself,  into  helium  and  niton. 

There  are  three  series  of  radioactive  elements:  the 
radium  series  beginning  with  uranium  and  ending  with 
lead ;  a  thorium  series  ending  with  bismuth ;  and  an  actinium 
series  ending  with  an  unknown,  inactive  element. 


276  ALKALI-EARTH  METALS 

It  is  proposed  to  use  mesothorium  in  radiolite  watches. 

The  " half-life  period"  of  an  element  is  the  time  required 
for  one-half  of  the  element  to  disintegrate. 

The  number  of  helium  atoms  shot  out  by  a  small  quantity 
of  radium  has  been  counted  and  by  this  method  the  number 
of  molecules  in  a  cubic  centimeter  of  a  gas  has  been 
determined. 

EXERCISES 

1.  How  much  limestone  will  be  required  to  give  a  ton  of  lime  ? 

2.  How  much  s^ked  lime  will  a  ton  of  lime  give? 

3.  How  much  barium  peroxide  will  be  required  to  give  one  liter 
of  a  3  per  cent  solution  of  hydrogen  peroxide? 

4.  The  sulfur  in  1.25  grams  of  coal  was  oxidized  to  a  sulfate 
and  the   solution   was   precipitated  with  a  solution  of  barium 
chloride.     The   barium   sulfate   obtained   weighed   0.183   gram. 
What  per  cent  of  sulfur  did  the  coal  contain? 


CHAPTER  XXIV 

METALLURGY    AND    THE    PREPARATION    OF    COM- 
POUNDS OF  THE    METALS 

Metallurgy. — By  metallurgy  is  meant  the  preparation 
of  a  metallic  element  in  the  free  state,  generally  for  practical 
use.  As  the  human  race  emerged  from  savagery  the  first 
metals  to  be  used  were  copper,  silver  and  gold,  all  of  which 
are  found  free  in  nature.  Tin  was  added  to  the  list  in 
comparatively  early  times  because  it  can  easily  be  reduced 
from  the  oxide,  which  is  the  most  common  ore.  Bronzes 
prepared  by  the  reduction  of  mixed  ores  were  also  used.  It 
was  not  till  a  comparatively  high  state  of  civilization  was 
reached,  though  still  in  prehistoric  times,  that  men  learned 
to  reduce  iron  from  its  ores  in  simple  furnaces  somewhat 
like  a  blacksmith's  forge. 

Iron  was  a  comparatively  rare  and  expensive  metal 
through  the  middle  ages.  About  1500,  however,  the  blast 
furnace  was  invented  for  the  reduction  of  iron  ores.  Coal 
was  first  used  in  blast  furnaces  in  1735  (see  p.  324).  The 
rapid  reduction  of  iron  on  a  large  scale  has  now  made  iron 
the  cheapest  of  all  the  metals  and  in  its  various  forms  it 
is  of  greater  importance  than  all  the  other  metals  combined. 

Reduction  of  Ores  of  the  Common  Metals. — The  most 
common  and  important  metals  are  iron,  copper,  zinc,  lead, 
tin  and  antimony.  All  these  are  reduced  to  the  metallic 
state  from  the  oxide  by  one  of  three  methods : 

1.  The  oxides  of  iron,  zinc  and  tin  are  reduced  by  fuels 
.containing  carbon 

2.  Sulficles  of  zinc  and  antimony  are  oxidized  to  sulfur 

277 


278  METALLURGY 

dioxide  and  the  oxide  of  the  metal  by  heating  them  in  the 
air  and  the  oxide  is  then  reduced  by  a  fuel  as  in  the  first 
method. 

3.  Sulfides  of  copper  and  lead  are  partially  oxidized  by 
roasting  in  the  air  and  the  mixture  of  oxide  or  sulfate  with 
some  of  the  original  sulfide  is  then  heated.  The  two  com- 
pounds reduce  each  other  and  sulfur  dioxide  escapes. 

Electrolytic  Methods. — Sir  Humphrey  Davy  discovered 
how  to  prepare  potassium  and  sodium  by  means  of  an 
electric  current  a  little  more  than  100  years  ago,  but  it  was 
not  till  near  the  close  of  the  nineteenth  century  that  methods 
of  producing  electric  currents  on  a  large  scale  by  means  of 
dynamos  were  developed.  Since  that  time  electricity  has 
been  extensively  used  in  the  refining  of  copper  and  in  the 
manufacture  of  aluminium  and  sodium  and  magnesium. 
The  properties  of  copper  are  such  that  it  may  easily  be  re- 
duced from  the  aqueous  solution  of  its  salts,  but  in  the 
commercial  production  of  aluminium  and  sodium  water  is 
excluded. 

Reduction  by  Means  of  Aluminium. — Still  more  recently, 
the  preparation  of  aluminium  by  electrolysis  has  made  it 
possible  to  use  the  metal  for  the  preparation  of  other  metals. 
The  use  of  a  mixture  of  ferric  oxide,  Fe2O3,  with  aluminium 
to  obtain  a  high  temperature  has  been  described  as  Gold- 
schmidt's  thermite  process.  If  chromic  oxide,  C^Os,  is 
used  in  place  of  ferric  oxide,  metallic  chromium  is  produced 
and  the  same  method  may  be  used  for  the  preparation  of 
other  metals. 

Preparation  of  Compounds  of  the  Metals. — Nearly  all 
compounds  of  the  metals  are  salts.  With  a  very  few  ex- 
ceptions these  are  prepared  by  reversible  reactions  carried 
out  in  solutions.  These  reactions  proceed  chiefly  in  one 
direction  because  of  one  of  the  three  following  conditions: 

1.  A  Volatile  Compound  is  Formed. — Because  the  vola- 
tile compound  escapes  from  the  mixture,  the  equilibrium 


PRECIPITATION  279 

of  the  reversible  reaction  is  constantly  shifted  toward  its 
formation : 

NaCl  +  H2SO4  *±  NaHSO4  +  HC1 

Na2CO3  +  2HC1  <=±  2NaCl  +  H2C03 

H2C03  <=*  H20  +  C02 

2.  A    Difficultly    Soluble    Compound  is  Formed. — The 

separation  of  such  a  compound  as  a  precipitate  shifts  the 
equilibrium  toward  its  formation: 

BaCl2  +  Na2S04  <=±  2NaCl  +  BaSO4 
NaHS04  +  HC1  <=±  H2SO4  +  NaCl 
NaN03  +  KC1  *±  KN03  +  NaCl 

In  the  first  illustration  the  barium  sulfate  is  so  nearly 
insoluble  that  the  reaction  is  practically  complete  and  it  is 
impossible  for  a  solution  containing  sulfate  ions  to  contain 
more  than  a  very  minute  trace  of  barium  ions. 

In  the  second  illustration  sodium  chloride  will  separate 
only  when  water  is  present  to  retain  the  hydrochloric  acid 
in  solution  and  when  the  amounts  of  sodium  and  chlorine 
are  more  than  sufficient  to  saturate  the  solution  with  sodium 
chloride.  In  both  this  and  the  third  case  it  is  evident  that 
salts  which  are  quite  soluble,  as  well  as  those  which  are 
ordinarily  spoken  of  as  insoluble,  may  be  obtained  by  reac- 
tions of  this  type.  Potassium  nitrate,  KNO3,  is  very  easily 
soluble  in  hot  water  while  common  salt,  NaCl,  is  scarcely 
more  soluble  in  hot  than  in  cold  water.  When  a  mixture  of 
sodium  nitrate  and  potassium  chloride  is  treated  with  a 
small  amount  of  hot  water,  sodium  chloride,  the  least 
soluble  of  the  four  salts,  remains  undissolved  and  by  pouring 
off  and  cooling  the  hot  solution  potassium  nitrate  is  obtained. 
Saltpeter  for  the  manufacture  of  gunpowder  is  obtained  in 
this  way. 

3.  A  Compound  Which  Ionizes  to  only  a  Slight  Extent  is 
Formed. — The  most  common  case  of  this  class  of  reactions 


280  METALLURGY 

is  the  formation  of  water  by  the  union  of  hydrogen,  H+, 
and  hydroxide,  OH~,  ions.  The  ionization  of  water: 

HOH  <=>  H+  +  OH- 

takes  place  to  such  a  trifling  extent  that  it  is  impossible  to 
haveTmore  than  a  very  few  hydrogen  ions  in  a  solution 
containing  hydroxide  ions  or  more  than  a  very  few  hydrox- 
ide ions  in  a  solution  containing  hydrogen  ions: 

NaOH  +  HN03  +±  NaN03  +  HOH 

Solubility  of  Salts. — Salts  vary  greatly  in  solubility  and 
no  satisfactory  reasons  can  be  given  why  some  salts  are 
soluble  while  others  are  insoluble.  In  spite  of  some  excep- 
tions the  following  general  rules  are  useful: 

1.  Practically   all   salts   of   the   alkali   metals,   sodium, 
potassium  and  ammonium,  are  soluble  in  water. 

2.  Nearly  all  salts  of  the  strong,  highly  ionized,  monobasic 
and   bibasic   acids   are   soluble.     This  includes   chlorides, 
bromides,    iodides,    fluorides,    chlorates    and  perchlorates, 
sulfites  and  sulfates,  nitrites  and  nitrates. 

3.  Normal  salts  of  the   weak,   slightly  ionized,  bibasic 
acids,  carbonic,  H2CO3,  silicic,  H2SiO3,  and  hydrosulfuric, 
H2S,  and  of  the  tribasic  acids,  phosphoric,  H3P04,  arsenious, 
H3AsO3,  arsenic,  HsAsO^  and  boric,  H3B03,  are  insoluble 
with  the  exception  of  the  salts  of  the  alkalies. 

The  following  exceptions  and  modifications  of  these 
rules  are  given  for  reference  but  should  be  learned  rather 
by  experience  in  the  laboratory  than  by  memorizing  them. 

Potassium  and  ammonium  chloroplatinates,  K2PtCl6 
and  (NH4)2PtCl6,  and  potassium  perchlorate,  KC104,  are 
only  slightly  soluble. 

The  following  salts  of  monobasic  and  bibasic  acids  are 
nearly  insoluble:  The  chlorides,  bromides  and  iodides  of  sil- 
ver, cuprous  copper,  mercurous  mercury  and  lead,  AgCl, 
AgBr,  Agl,  CuCl,  Cul,  HgCl,  Hgl,  PbCl2  (slightly  soluble), 


SUMMARY.     METALLURGY  281 

PbBr2,  PbI2,  mercuric  iodide,  HgI2,  calcium  fluoride,  CaF2, 
barium  sulfite,  BaSO3,  and  calcium,  strontium,  barium  and 
radium  sulfates,  CaS04  (slightly  soluble),  SrS04,  BaS04, 
RaSO4. 

Sulfides  of  the  alkali  metals,  sodium,  potassium  and 
ammonium,  and  of  the  alkali-earth  metals,  calcium,  barium 
and  strontium,  are  hydrolyzed  by  water  to  a  hydrosulfide 
and  hydroxide: 

Na2S  +  HOH  *±  NaHS  +  NaOH 
2CaS  +  2HOH  <=»  Ca(SH)2  +  Ca(OH)2 

The  sulfides  of  aluminium  and  chromium  are  hydrolyzed 
to  hydroxides  and  hydrogen  sulfide : 

A12S3  +  6HOH  «=±  2A1(OH)3  +  3H2S 

The  sulfides  of  all  the  other  common  metals,  except 
magnesium,  are  insoluble  in  water. 

SUMMARY 

The  most  important  ores  of  the  common  metals  are  oxides 
or  sulfides. 

Common  metals  are  obtained  by  reducing  the  oxide,  in 
nearly  all  cases.  This  may  be  preceded  by  roasting  the 
sulfide. 

Some  metals  are  obtained  and  others  are  purified  by 
electrolysis. 

A  few  of  the  less  common  metals  are  prepared  by  re- 
ducing the  oxide  with  aluminium. 

Nearly  all  reactions  for  the  preparation  of  salts  are 
reversible  and  depend  on  the  escape  of  a  volatile  compound, 
the  precipitation  of  an  insoluble  compound,  or  the  formation 
of  a  slightly  ionized  compound. 

Salts  of  the  alkali  metals  and  of  strong  monobasic  and 
bibasic  acids  are  usually  soluble,  but  almost  all  salts  of 
weak  bibasic  acids  and  of  tribasic  acids  are  insoluble. 


282  METALLURGY 

EXERCISES 

Write  the  equations  for  the  following  processes: 

1.  The  reduction  of  ferric  oxide  by  carbon  monoxide. 

2.  The  roasting  of  antimony  sulfide  to  the  oxide  and  the  reduc- 
tion of  the  latter  to  metallic  antimony. 

3.  The  roasting  of  galena  to  lead  oxide  and  lead  sulfate  and  the 
reduction  of  each  by  heating  it  with  lead  sulfide. 

4.  The  reduction  of  chromic  oxide  by  aluminium. 


CHAPTER  XXV 

GROUP    II:   SECOND    DIVISION;    MAGNESIUM,    ZINC, 
CADMIUM  AND  MERCURY 

Characteristics  of  the  Metals  of  the  Second  Division. — 

While  magnesium  resembles  calcium  both  in  its  chemical 
and  physical  .properties  and  in  its  occurrence  in  nature,  zinc, 
cadmium  and  mercury  differ  in  increasing  degree  from  stron- 
tium, barium  and  radium  (see  p.  260).  The  hydroxides 
of  magnesium,  zinc  and  cadmium  are  insoluble  and  mercury 
forms  no  hydroxide,  while  the  hydroxides  of  strontium  and 
barium  are  more  soluble  than  that  of  calcium  and  are  strong 
bases. 

Magnesium.  Occurrence. — The  occurrence  of  magne- 
sium with  calcium  in  the  mineral  dolomite,  MgCO3.CaCO3, 
has  been  referred  to.  The  pure  carbonate,  MgCO3,  is 
known  as  the  mineral  magnesite.  Magnesium  sulfate, 
MgSO4.7H2O,  is  easily  soluble  and  is  found  in  some  well- 
known  mineral  waters,  especially  in  Hunyadi  water,  to 
which  it  gives  part  of  its  medicinal  value. 

Metallic  Magnesium  is  a  very  light  metal,  having  a 
specific  gravity  of  only  1.8.  It  resembles  zinc  in  its  appear- 
ance and  in  some  of  its  properties,  but  is  much  more 
active.  It  decomposes  boiling  water  with  the  evolution  of 
hydrogen.  In  the  form  of  a  ribbon  it  can  be  easily  burnt 
in  air,  giving  a  very  intense  white  light.  In  the  form  of  a 
powder  mixed  with  potassium  chlorate,  which  yields  oxygen- 
readily,  it  is  used  in  flash-light  powders  for  photography. 
This  is  because  the  intense  white  light  given  when  it  burns 
contains  the  violet  rays  which  affect  the  salts  of  silver  used 
in  photographic  plates. 

283 


284    MAGNESIUM,  ZINC,  CADMIUM  AND  MERCURY 

Compounds  of  Magnesium. — Magnesium  forms  the  com- 
pounds to  be  expected  of  a  bivalent  metal  easily  soluble  in 
acids:  the  oxide,  MgO,  hydroxide,  Mg(OH)2,  chloride, 
MgCl2.6H2O,  and  sulfate,  MgSO4.7H2O. 

Zinc.  Occurrence,  Metallurgy. — Zinc  is  found  chiefly 
as  the  sulfide,  sphalerite,  ZnS,  but  the  carbonate,  ZnCOs, 
and  a  hydrous  silicate,  ZnSiO3.H2O,  are  also  found  in  suffi- 
cient amounts  to  be  available  ores. 

To  obtain  metallic  zinc  the  sulfide  is  roasted,  giving  sulfur 
dioxide,  SO2,  and  zinc  oxide,  ZnO.  The  sulfur  dioxide  is 
now  often  used  for  the  manufacture  of  sulfuric  acid.  This 
is  done  partly  to  prevent  the  escape  of  sulfur  dioxide  which 
would  injure  the  vegetation  in  the  neighborhood  of  the 
works. 

The  zinc  oxide  is  mixed  with  coal  and  heated  in  an  earth- 
enware retort  which  has  an  earthenware  receiver,  in  which 
the  zinc  that  distils  from  the  retort  collects. 

Properties  of  Zinc.  Alloys. — Zinc  is  more  electropositive 
than  iron  (p.  243).  In  spite  of  this,  it  is  much  less  affected 
than  iron  by  the  combined  action  of  air  and  water.  This 
is  due  in  part  to  the  fact  that  the  ferric  oxide  and  hydroxide, 
which  we  call  iron  rust,  does  not  adhere  closely  to  the  iron 
and  the  difference  of  electrical  potential  between  the  rust 
and  the  iron  hastens  the  action  of  air  and  water.  The 
oxide  or  hydroxide  formed  on  the  surface  of  zinc,  on  the 
other  hand,  adheres  closely  and  forms  a  coating  which  pro- 
tects the  metal  from  further  action.  This  property  of 
zinc,  together  with  the  fact  that  it  is  more  electropositive 
than  iron  and  so  protects  iron  which  is  in  contact  with  it 
from  rusting,  renders  zinc  more  effective  than  tin  as  a  pro- 
tective coating  for  iron.  It  must  be  remembered,  however, 
that  zinc  dissolves,  even  in  weak  acids,  and  that  zinc  com- 
pounds are  poisonous.  This  makes  iron  covered  with  zinc, 
commonly  called  "galvanized  iron,"  unsuitable  for  kitchen 
utensils. 


ZINC.     MERCURY  285 

Zinc  is  used  as  the  electropositive  metal,  which  passes 
into  solution,  in  all  kinds  of  primary  electrical  batteries, 
including  the  well-known  "dry"  cells. 

Brass  is  an  alloy  of  approximately  two  parts  of  copper 
with  one  of  zinc. 

Zinc  Sulfide,  ZnS,  is  found  in  nature  as  the  mineral  sphal- 
erite. It  may  be  prepared  as  a  white  precipitate  when 
hydrogen  sulfide  is  passed  into  a  neutral  or  weakly  acid 
solution  of  a  zinc  salt.  It  is  the  only  white  metallic  sulfide 
which  can  be  prepared  in  this  way. 

Zinc  Oxide,  ZnO,  is  prepared  by  burning  metallic 
zinc.  It  is  a  white  powder  which  forms  an  excellent  white 
paint  when  mixed  with  linseed  oil.  White  lead  and  paints 
which  contain  it  are  blackened  by  hydrogen  sulfide  because 
lead  sulfide,  PbS,  is  black.  Since  zinc  sulfide  is  white, 
paints  containing  zinc  oxide  are  not  blackened  in  the  same 
way  and  such  paints  are  more  suitable  than  paints  contain- 
ing white  lead  for  use  in  laboratories  and  in  situations  where 
the  paint  is  exposed  to  sewage  gases. 

Zinc  Sulfate,  ZnSO4.7H2O.  Vitriols. — The  name  "  vit- 
riol "  seems  to  have  been  first  used  for  ferrous  sulfate  or  green 
vitriol,  FeSO4.7H2O,  and  later  for  sulfuric  acid,  which  was 
called  "oil  of  vitriol"  because  it  was  prepared  by  the  dis- 
tillation of  a  mixture  of  sulfates  of  iron.  After  that  zinc 
sulfate  was  called  white  vitriol  and  copper  sulfate,  CuSO-i.- 
5H2O,  blue  vitriol. 

Cadmium  closely  resembles  zinc  in  most  of  its  properties. 
Its  sulfide,  CdS,  is  yellow.  The  metal  is  used  in  some  of 
the  fusible  alloys. 

Mercury.  Metallurgy,  Properties. — Mercury  is  found 
to  a  limited  extent  in  the  free  state  in  nature  but  occurs 
chiefly  as  the  sulfide,  HgS,  in  the  red  mineral  cinnabar. 
The  metal  is  prepared  from  this,  either  by  roasting  it  to- 
sulfur  dioxide  and  metallic  mercury,  or  by  mixing  it  with 
iron  and  distilling  away  the  mercury. 


286    MAGNESIUM,  ZINC,  CADMIUM  AND  MERCURY 

Mercury  is  a  heavy,  mobile  liquid.  It  melts  at  nearly  40° 
below  zero,  the  Fahrenheit  and  Centigrade  thermometers 
coming  together  at  that  point.  It  boils  at  about  360°. 
The  specific  gravity  is  13.6.  Because  it  is  a  liquid  through 
such  a  convenient  range  of  temperatures  and  because  its 
rate  of  expansion  is  very  uniform  when  it  is  heated,  it  is 
used  for  thermometers.  Because  of  its  specific  gravity, 
such  that  a  column  760  mm.  or  30  inches  in  height 
will  balance  the  pressure  of  the  atmosphere}  it  is  used  for 
barometers. 

Mercury  alloys  easily  with  gold,  silver  and  many  other 
metals.  These  alloys  are  called  amalgams.  If  the  amount 
of  the  foreign  metal  is  small,  the  metal  dissolves  in  the 
mercury  and  the  amalgam  remains  liquid,  but  in  most 
cases  only  a  small  per  cent  of  the  foreign  metal  is  required 
to  give  a  solid  amalgam.  The  property  of  amalgamating 
with  gold  and  silver  is  used  in  separating  these  metals 
from  the  large  quantities  of  other  minerals  with  which  they 
are  usually  mixed. 

Valence  of  Mercury. — All  of  the  metals  of  both  divisions 
of  Group  II  except  mercury  form  exclusively  compounds 
in  which  the  metals  are  bivalent,  as  for  instance  in  the 
chlorides  CaCl2  and  CdCl2  and  the  sulfates  BaSO4  and 
ZnSO4.  Mercury,  on  the  other  hand,  forms  mercurous 
compounds,  such  as  Hg2O  and  HgCl  (or  Hg2Cl2)  in  which  it 
appears  univalent,  as  well  as  mercuric  compounds,  such  as 
HgO,  HgCl2  and  HgSO4,  in  which  it  is  bivalent. 

Mercuric  Oxide,  HgO. — In  the  second  chapter  an  experi- 
ment of  Lavoisier  was  described  in  which  he  heated  some 
mercury  in  contact  with  a  limited  amount  of  air  till  the 
oxygen  had  been  converted  into  the  red  oxide  of  mercury. 
He  then  recovered  the  oxygen  by  heating  the  oxide  to  a 
higher  temperature,  decomposing  it  into  mercury  and  oxy- 
gen. Priestly  had  previously  obtained  oxygen  in  England 
by  heating  this  same  oxide.  The  historical  interest  at- 


CALOMEL.     MERCURIC  FULMINATE  287 

tached  to  these  experiments  has  made  mercuric  oxide  a 
common  substance  in  chemical  laboratories. 

Mercurous  Chloride,  Hg2Cl2,  or  Calomel. — Mercurous 
chloride  is  prepared  by  subliming  a  mixture  of  mercuric 
chloride,  HgCl2,  and  mercury: 

HgCl2  +  Hg  =  Hg2Cl2 

Calomel  is  extensively  used  as  a  medicine.  A  widespread 
prejudice  against  its  use  arose  from  its  administration 
in  large  doses,  sometimes  causing  salivation  and  other 
serious  injuries.  It  is  now  usually  given  in  very  small  doses 
and  is  mixed  with  sodium  bicarbonate,  NaHCOs,  to  neutral- 
ize the  acid  of  the  gastric  juice  and  render  the  calomel  less 
soluble. 

Calomel  is  practically  insoluble  in  dilute  acids  and  is 
precipitated  when  hydrochloric  acid  or  a  chloride  is  added 
to  a  solution  of  a  soluble  mercurous  salt. 

Mercuric  Chloride,  HgCl2,  or  Corrosive  Sublimate. — 
Mercuric  chloride  is  moderately  soluble  in  water  and  easily 
soluble  in  alcohol.  The  best  antidote  is  the  white  of  an  egg, 
with  which  it  forms  an  insoluble  compound,  but  this  ant\- 
dote,  to  be  effective,  must  be  administered  very  promptly. 

Mercuric  chloride  is  a  very  powerful  antiseptic  and  is 
extensively  used  in  antiseptic  surgery. 

Mercuric  Fulminate,  Hg(CNO)2,  a  salt  of  fulminic  acid, 
HCNO,  is  used  as  a  primer  in  cartridges  and  in  the  per- 
cussion caps  used  to  explode  dynamite.  It  explodes 
easily  with  a  blow,  decomposing  into  carbon  monoxide, 
CO,  nitrogen,  N2,  and  mercury,  all  of  which  are  gases  at  the 
temperature  of  the  explosion. 

SUMMARY 

The  metals  of  the  second  division  of  Group  II,  with  the 
exception  of  mercury,  form  insoluble  hydroxides.  Mercury 
forms  no  stable  hydroxide.  Those  with  the  high  atomic 


288    MAGNESIUM,  ZINC,  CADMIUM  AND  MERCURY 

weights  differ  very  markedly  from  the  metals  of  the  first 
division. 

Magnesium  occurs  as  dolomite  and  magnesite. 

Metallic  magnesium  resembles  zinc  but  is  much  lighter 
and  more  active.  It  is  used  in  flash-light  powders. 

Magnesium  is  bivalent.  The  oxide,  hydroxide,  chloride 
and  sulfate  are  common  compounds. 

Zinc  occurs  as  sphalerite  and  as  the  carbonate  and 
silicate.  It  is  reduced  by  heating  the  oxide  with  carbon 
and  distilling  the  zinc. 

Zinc  sulfide  is  white. 

Zinc  oxide  is  used  as  a  white  paint,  which  is  not  darkened 
by  hydrogen  sulfide. 

Zinc  chloride  is  used  as  a  disinfectant. 

Zinc  sulfate  is  white  vitriol. 

Cadmium  is  used  in  fusible  alloys. 

Mercury  occurs  as  cinnabar.  It  is  obtained  by  roasting 
the  ore  or  by  mixing  it  with  iron  and  distilling. 

It  is  a  heavy  liquid,  used  in  barometers  and  thermometers. 
It  forms  amalgams  with  many  metals.  The  gold  and  silver 
amalgams  are  used  in  the  recovery  of  gold  and  silver  from 
their  ores. 

Mercury  is  bivalent  in  some  compounds,  univalent  in 
others. 

Mercurous  chloride  or  calomel  is  used  as  a  medicine. 
Mercuric  chloride,  or  corrosive  sublimate,  is  used  as  an 
antiseptic  in  surgery. 

Mercuric  fulminate  is  used  in  primers  and  fulminating  caps 
to  detonate  nitroglycerine  and  other  explosives. 

EXERCISES 

1.  Write  the  equations  for  the  metallurgy  of  zinc,  starring 
with  sphalerite. 

2.  Barium  sulfate  may  be  reduced  to  the  sulfide  by  heating 
it  with  carbon.     The  barium  sulfide  will  dissolve  in  water  and  the 
solution  is  made  to  react  with  a  solution  of  zinc  sulfate  to  produce 


EXERCISES.     MAGNESIUM,  ZINC,  MERCURY       289 

the  commercial  pigment  called  lithopone.     Write  the  equations 
and  explain  the  process. 

3.  What  is  the  difference  between  sublimation  and  distillation? 

4.  Mercury  is  converted  into  mercuric  sulfate  by  heating  with 
concentrated  sulfuric  acid,  a  part  of  the  acid  being  reduced  to 
sulfur  dioxide.     Corrosive  sublimate  is  prepared  by  subliming 
a  mixture  of  mercuric  sulfate  with  salt.     Write  the  equations. 


CHAPTER  XXVI 

GROUP  I:  FIRST  DIVISION;  ALKALI  METALS,  SODIUM, 
POTASSIUM  AND  AMMONIUM.     SPECTRUM  ANALYSIS 

Properties  of  the  Alkali  Metals. — The  halogens,  chlorine, 
bromine,  iodine  and  fluorine,  which  are  univalent  toward 
hydrogen,  stand  at  the  extreme  in  non -metallic  properties 
and  in  reactivity.  In  a  similar  manner  the  alkali  metals, 
lithium,  sodium,  potassium,  rubidium  and  caesium,  which  are 
univalent  toward  chlorine  and  other  halogens,  stand  at  the 
extreme  in  metallic  properties  and  in  reactivity.  Freshly 
cut  surfaces  of  these  metals,  or  surfaces  which  are  protected 
from  the  action  of  the  air  and  moisture  by  melting  in  an 
evacuated  tube,  have  a  bright  luster  resembling  that  of 
silver.  The  metals  react  violently  with  water,  liberating 
hydrogen  and  forming  hydroxides  which  are  strong  bases, 
easily  soluble  in  water. 

With  a  very  few  exceptions  the  salts  of  the  alkali  metals 
are  easily  soluble  in  water.  The  salts  are  highly  ionized  in 
solution  and  the  solutions  are  good  conductors  of  electricity. 

Sodium. — The  occurrence  of  sodium  as  common  salt, 
NaCl,  and  the  preparation  and  properties  of  metallic  sodium 
were  described  in  an  early  chapter  of  the  book  and  should  be 
reviewed  at  this  point.  A  number  of  additional  facts  will 
be  given  here. 

Besides  the  occurrence  as  sodium  chloride,  sodium  is  found 
in  very  many  of  the  natural  silicates  and  also  as  sodium 
nitrate  or  Chile  saltpeter,  NaNO3. 

Sodium  Chloride,  NaCl. — Salt  is  probably  the  most  com- 
mon mineral  constituent  of  our  food,  though  only  a  small 

290 


SODIUM  291 

quantity  is  required.  By  the  physiological  processes  of 
the  body  it  is  hydrolyzed,  yielding  hydrochloric  acid,  which 
is  an  essential  part  of  the  gastric  juice,  the  digestive  fluid 
of  the  stomach.  The  sodium,  on  the  other  hand,  is  changed 
to  sodium  bicarbonate,  NaHCO3,  or  disodium  phosphate, 
Na2HPO4,  which  are  important  constituents  of  the  blood 
and  of  the  alkaline  fluids  of  the  digestive  tract. 

Salt  is  the  source  for  the  manufacture  of  chlorine,  hydro- 
chloric acid  and  of  practically  all  important  compounds  of 
chlorine.  It  is  also  the  ultimate  source  for  the  manufacture 
of  sodium  hydroxide,  sodium  carbonate  and  all  important 
compounds  of  sodium  except  sodium  nitrate.  These  com- 
pounds include  soap  and  glass.  It  will  be  seen  from  these 
staternents  that  salt  fills  a  unique  and  fundamental  place 
in  our  chemical  industries. 

Sodium  Nitrate,  NaNO3. — For  several  generations  the 
Chile  saltpeter,  coming  from  a  limited  area  on  the  west  coast 
of  South  America,  has  furnished  nearly  all  of  the  nitric  acid 
and  saltpeter,  KNO3,  of  the  world.  Through  these  com- 
pounds the  explosives  used  in  blasting,  hunting  and  warfare 
have  been  made.  Nitric  acid  is  also  used  in  the  manufac- 
ture of  dyes  and  of  a  number  of  important  compounds  used 
in  medicine.  Sodium  nitrate  is  also  extensively  used  in 
fertilizers  applied  to  soils  deficient  in  compounds  of  nitro- 
gen. It  is  estimated  that  the  present  source  of  the  mineral 
will  be  exhausted  within  a  comparatively  few  years. 
Fortunately  methods  of  making  nitrates  and  nitric  acid 
from  the  nitrogen  of  the  air  have  been  discovered  and  there 
is  no  probability  of  any  serious  injury  either  to  our  manu- 
factures or  to  agriculture  when  the  beds  of  sodium  nitrate 
in  Chile  are  gone. 

Sodium  Carbonate,  Sal  Soda,  or  Washing  Soda, 
Na2CO3.10H2O.  Le  Blanc  Soda  Process. — During  the 
nineteenth  century  a  large  part  of  the  sodium  carbonate 
used  for  the  manufacture  of  soap  and  glass  was  manu- 


292  ALKALI  METALS 

factured  by  a  process  invented  by  Le  Blanc  just  before 
the  French  revolution.  The  process  consists  of  three 
operations : 

1.  Treatment  of  salt  with  sulfuric  acid  and  heating  the 
mixture  till  the  hydrochloric  acid  is  expelled: 

NaCl  +  H2S04     =   NaHSO4  +  HC1 
Nad  +  NaHSQ,  =  Na2SO4    +  HC1 

2.  Heating  the  sodium  sulfate  with  carbon  and  calcium 
carbonate,  to  reduce  the  sodium  sulfate  to  sodium  sulfide 
and  convert  the  latter  into  sodium  carbonate : 

Na2SO4  +    2C  '     =  Na2S       +  2CO2 
Na2S      +  CaCO3  =  Na2CO3  +  CaS 

3.  Separation  of  the  sodium  carbonate  from  the  calcium 
sulfide  by  dissolving  the  former  in  water.     The  calcium 
sulfide  is  nearly  insoluble  in  the  alkaline  solution  and  may  be 
separated  from  it.     The  sal  soda  may  then  be  obtained  by 
evaporating  and  cooling  the  clear  solution. 

Alkalinity  of  Sodium  Carbonate. — A  solution  of  sodium 
carbonate  turns  red  litmus  blue  and  reacts  alkaline  toward 
all  of  the  common  indicators.  Normal  sodium  salts  or 
potassium  salts  of  many  other  weak  acids  give  similar  alka- 
line solutions.  We  may  say,  superficially  and  with  some 
truth,  that  these  weak  acids  are  not  able  to  completely  neu- 
talize  strong  bases,  such  as  sodium  hydroxide  and  potassium 
hydroxide.  The  phenomena  may  be  more  fully  and  accu- 
rately explained  by  a  consideration  of  the  equilibria  rep- 
resented in  the  following  equations: 

HOH      <=»  H+     +  OH- 
Na2CO3  <=±  Na+  +  Na+  +  C03- 
C03=  +  H+  <=±  HC03- 

combining: 

Na2CO3  +  HOH  <=±  Na+  +  Na+  +  OH~  +  HCO3~ 


SOLVAY  PROCESS  293 

In  pure  water  and  in  any  truly  neutral  solution  the  numbers 
of  the  hydrogen,  H+,  and  hydroxide,  OH~,  ions  are  equal. 
When  sodium  carbonate  is  dissolved  in  the  water  the  car- 
bonate ions,  CO3=,  combine  with  the  hydrogen  ions  of  the 
water  to  form  hydrocarbonate  ions,  HCO3~,  first,  because 
the  hydrocarbonate  ions  separate  to  only  a  very  trifling 
extent  into  hydrogen,  H+,  and  carbonate,  CO3=,  ions,  and 
second,  because  even  this  slight  degree  of  separation  is 
repressed  by  the  large  number  of  carbonate  ions  which  are 
present. 

Similar  explanations  may  be  given  in  the  case  of  other 
salts  of  weak  acids. 

Sodium  Bicarbonate  or  Baking  Soda,  NaHCO3 ;  Ammonia- 
Soda  or  Solvay  Process. — About  1860  Ernst  Solvay  of  Brus- 
sels, Belgium,  began  the  commercial  development  of  a  process 
which  had  been  discovered  twenty  years  before.  In  this 
process  ammonia  is  added  to  a  salt  brine  and  carbon  dioxide 
is  passed  into  the  mixture: 

NH3  +  H2O    +  CO2  =  NH4HCO3 

Ammonium  bicarbonate 

NH4HCO3  +  NaCl  <=>  NH4C1  +  NaHCO3 

Sodium  bicarbonate 

The  process  depends,  of  course,  on  the  fact  that  the  so- 
dium bicarbonate,  NaHCO3,  is  the  least  soluble  of  the  four 
compounds  given  in  the  second  equation. 

Ammonia  and  ammonium  chloride  are  much  more  valu- 
able than  the  equivalent  amounts  of  sodium  bicarbonate. 
The  commercial  success  of  the  process,  therefore,  depends 
on  recovering  the  ammonia  and  using  it  again  indefinitely. 
This  is  brought  about  by  treating  the  solution  from  which 
the  sodium  bicarbonate  has  been  separated  with  slaked 
lime: 

2NH4C1  +  Ca(OH)2  =  CaCl2  +  2NH3  +  2H2O 

The  ammonia  is  volatile  and  can  be  easily  separated  from 
the  calcium  chloride  by  distillation.  The  ammonia-soda 


294 


ALKALI  METALS 


process  has  been  found  to  be  more  economical  than  the  Le 
Blanc  soda  process  and  by  the  close  of  the  nineteenth 
century  the  competition  had  almost  completely  driven  out 
the  latter  method  of  manufacture. 

It  will  be  noticed  that  the  principles  of  solubility  and  vola- 
tility (p.  278)  are  used  in  both  the  Le  Blanc  and  the  Solvay 
processes. 

The  use  of  sodium  bicarbonate  in  baking  powders  has 
been  referred  to  in  a  previous  chapter.  It  is  easily  decom- 
posed by  heat  into  sodium  carbonate,  water  and  carbon 
dioxide. 

Sodium  Hydroxide,  NaOH. — As  long  as  sodium  car- 
bonate was  prepared  by  the  Le  Blanc  process  or  by  the 
ammonia-soda  process,  sodium  hydroxide  for  the  manu- 
facture of  soap  and  for  other  uses  was  prepared  by  adding 
slaked  lime,  Ca(OH)2,  to  a  solution  of  sodium  carbonate. 
Calcium  carbonate,  CaC03,  is  less  soluble  than  calcium 
hydroxide  and  it  is  also  the  least  soluble  of  the  four 
substances  of  the  equation : 

Na2CO3  +  Ca(OH)2  =  CaC03  +  2NaOH 

Electrolytic  Sodium  Hydroxide. — The  development  of 
cheap  sources  of  electrical  power  has  lead  to  the  invention 


4- 


FIG.  48. 


of  many  different  forms  of  apparatus  for  the  electrolysis 
of  a  solution  of  salt,  NaCl,  in  such  a  manner  as  to  produce 
chlorine  and  sodium  hydroxide,  both  of  which  are  valu- 
able and  extensively  used.  One  of  the  best  of  these  is  the 
Castner-Kellner  apparatus  shown  in  Fig.  48.  It  consists 
of  a  slate  box  divided  into  three  compartments  by  two  par- 


SODIUM  HYDROXIDE  295 

titions,  which  fit  loosely  into  grooves  in  the  bottom  of  the 
box.  Mercury  placed  on  the  bottom  seals  these,  giving 
a  continuous  metallic  layer  for  the  three  compartments 
but  preventing  a  dilute  solution  of  sodium  hydroxide  placed 
in  the  central  compartment  from  mixing  with  the  brine 
in  the  two  side  compartments.  Graphite  anodes  are  used 
in  the  two  side  compartments  and  an  iron  cathode  in  the 
central  one.  Chlorine  is  evolved  from  the  anodes  and  is 
collected  and  used  for  the  manufacture  of  bleaching  powder 
or  for  some  other  purpose.  The  mercury  in  the  two  side 
compartments  is  negative  as  compared  with  the  graphite 
anodes  and  the  sodium  liberated  at  its  surface  combines 
with  it  to  form  a  liquid  sodium  amalgam.  By  a  slight  tilt- 
ing motion  the  amalgam  is  caused  to  flow  alternately  to 
one  side  and  the  other  and  is  brought  into  the  central  com- 
partment. Here  it  is  positive  with  reference  to  the  more 
negative  cathode  and  the  hydroxide  ions  brought  to  its 
surface  combine  with  the  sodium  of  the  amalgam  to  form 
sodium  hydroxide,  while  the  hydrogen  ions  of  the  water 
are  discharged  as  free  hydrogen  at .  the  surface  of  the  iron 
cathode.  The  hydroxide  solution  is  kept  at  a  constant 
concentration  by  introducing  water  at  one  side  and  remov- 
ing some  of  the  solution  at  the  other.  Salt  is  added  to  the 
side  compartments  from  time  to  time. 

Dry  sodium  hydroxide  is  obtained  from  the  solution  by 
evaporation  but  it  must  be  heated  to  nearly  a  red  heat  to 
remove  the  last  of  the  water.  It  melts  at  that  temperature 
and  may  be  cast  into  sticks.  The  pure  hydroxide  is  a  white 
solid,  which  absorbs  water  greedily  from  the  air  and  deli- 
quesces. The  solution  absorbs  carbon  dioxide,  however, 
and  the  solution  of  sodium  carbonate  formed  will  evaporate, 
leaving  a  white  residue. 

Sodium  Oxide,  Na2O,  may  be  prepared  by  heating  sodium 
hydroxide  with  metallic  sodium.  It  is  very  rarely  prepared 
or  used. 


296  ALKALI  METALS 

Sodium  Peroxide,  Na2O2,  is  prepared  by  heating  metallic 
sodium  in  a  current  of  air.  The  process  is  carried  out  in 
shallow  aluminium  trays.  It  is  now  extensively  manu- 
factured and  is  used  for  the  helmets  used  in  the  mine  rescue 
work  and  to  some  extent  for  regenerating  the  air  of  subma- 
rine boats.  In  contact  with  carbon  dioxide  and  moisture, 
sodium  carbonate  and  oxygen  are  formed. 

With  cold,  dilute  acids,  sodium  peroxide  gives  hydrogen 
peroxide,  H2O2,  which  is  used  to  bleach  silk,  wool,  hair  and 
other  substances,  which  would  be  injured  by  chlorine. 

Sodium  Sulfite,  Na2SOs.H2O,  is  prepared  by  burning 
sulfur  and  passing  the  sulfur  dioxide  formed  through  a 
solution  of  sodium  carbonate.  It  is  used  as  a  reducing 
agent  in  photographic  developers. 

Sodium  Thiosulfate,  Na2S2O3.sH2O,  is  prepared  by  dis- 
solving sulfur  in  a  solution  of  sodium  sulfite,  Na2SO3.  The 
name  is  given  because  it  may  be  considered  as  sodium  sulfate, 
Na2S(>4,  in  which  one  atom  of  oxygen  has  been  replaced 
by  an  atom  of  sulfur.  The  prefix  "thio"  is  derived  from  a 
Greek  word  meaning  sulfur. 

Sodium  thiosulfate  is  usually  called  "  sodium  hyposulfite" 
by  pharmacists  and  photographers.  It  is  used  in  "fixing" 
photographs  by  dissolving  those  portions  of  the  silver 
chloride  or  bromide  which  have  not  been  affected  by  the 
light. 

Sodium  Silicate,  or  Soluble  Glass,  Na2SiO3,  is  made  by 
melting  a  mixture  of  sand  and  sodium  carbonate.  It  dis- 
solves in  water  but  is  hydrolyzed  to  sodium  hydroxide  and 
colloidal  silicic  acid,  giving  an  alkaline  solution.  The  solu- 
tion is  used  to  fireproof  wood  and  fabrics,  covering  them 
with  a  thin,  glassy  coating,  which  renders  them  less  in- 
flammable. The  solution  (about  10  per  cent)  is  also  used 
in  preserving  eggs.  A  thin  layer,  impervious  to  the  bac- 
teria which  would  cause  the  decay  of  the  eggs,  is  formed  on 
their  surface. 


POTASSIUM.     SOFT  SOAP  297 

Borax,  Na2B4O7.10H2O,  was  described  under  boron 
(p.  253). 

Potassium.  Occurrence. — The  decomposition  and  dis- 
integration of  natural  silicates  during  the  geologic  ages, 
with  the  formation  of  soils,  clays  and  shales,  has  been  spoken 
of  in  connection  with  the  occurrence  of  aluminium.  All 
fertile  soils  still  contain  portions  of  the  original  minerals 
and  these  still  yield  potassium  or  sodium,  or  compounds  of 
these  elements,  in  such  a  form  that  the  potassium,  especially, 
is  taken  up  by  trees  and  plants  growing  in  the  soil.  Soils 
from  which  the  potassium  has  been  removed  or  depleted 
by  the  growth  of  crops  for  a  series  of  years  may  become  in- 
fertile for  that  reason.  The  three  elements  in  which  soils 
are  most  likely  to  become  deficient  are  potassium,  phos- 
phorus and  nitrogen. 

Wood  Ashes.  Soft  Soap. — When  wood  is  burned,  the 
larger  part  of  the  potassium  which  it  contains  is  left  in  the 
form  of  potassium  carbonate,  which  can  be  readily  obtained 
from  the  ashes  by  extracting  them  with  water  and  evapo- 
rating the  solution.  In  wooded  countries  this  process  is 
still  sometimes  carried  out  in  a  simple  way  by  taking  a  bar- 
rel, boring  a  hole  in  the  bottom,  covering  this  with  hay 
or  shavings  to  act  as  a  filter  and  filling  it  with  wood  ashes. 
The  ashes  are  then  "leached"  by  pouring  water  on  the  top 
and  allowing  it  to  run  through  the  mass.  Sometimes  lime 
is  put  in  the  bottom  of  the  barrel  to  change  the  potassium 
carbonate,  K2CO3,  to  potassium  hydroxide,  KOH.  The 
"lye"  obtained  in  this  way  is  concentrated  by  boiling  it 
down  in  an  iron  pot,  if  necessary,  and  soft  soap  is  made  from 
it  by  boiling  it  with  the  refuse  grease  of  the  household. 

The  process  of  making  soap  which  has  been  described  was 
common  in  America  a  few  generations  ago  but  has  now  al- 
most disappeared,  partly  with  the  destruction  of  our  forests 
and  the  passing  of  wood  as  a  fuel,  partly  with  the  introduc- 
tion of  cheap  methods  for  the  manufacture  of  sodium  car- 


298  ALKALI  METALS 

bonate  and  sodium  hydroxide  and  the  substitution  of  the 
hard  soaps  made  from  these  for  the  soft  soaps  formerly  used. 

Potash  Salts  from  Germany. — Compounds  of  potassium 
are  still  required  in  large  quantities  for  the  manufacture  of 
saltpeter,  KNO3,  and  for  fertilizers  to  be  used  on  soils 
which  are  deficient  in  the  element.  At  Stassfurt,  in  Ger- 
many, there  are  deposits  of  minerals  which  contain  potas- 
sium chloride  and  for  some  time  past  this  has  been  the 
cheapest  source  of  potassium  compounds  for  the  whole 
world.  During  the  European  war  this  supply  has  been  cut 
off  and  a  very  careful  search  has  been  made  for  potassium 
salts  which  may  be  obtained  from  other  sources.  Among  the 
most  promising  of  these  may  be  mentioned  natural  silicates 
which  can  be  decomposed  by  processes  which  have  not  yet 
proved  to  be  commercially  successful.  Potassium  salts 
are  also  found  in  seaweeds  of  the  Pacific  Coast  and  serious 
attempts  are  being  made  to  utilize  these.  Considerable 
quantities  of  potassium  compounds  have  been  obtained 
from  the  water  of  Searle's  Lake,  California,  from  some  Ne- 
braska lakes,  and  also  as  a  by-product  in  the  manufacture  of 
cement.  It  is  too  soon  to  know  whether  the  potassium 
compounds  from  any  other  source  can  compete  with  the 
German  supply  after  the  war. 

Metallic  Potassium. — In  1807  Sir  Humphrey  Davy 
awakened  very  great  interest  in  the  scientific  world  by 
applying  the  current  from  a  powerful  electric  battery  to 
moist  potassium  hydroxide  and  preparing  in  this  way 
minute  globules  of  metallic  potassium.  He  also  tried  the 
experiment  with  moist  sodium  hydroxide  and  obtained 
metallic  sodium.  Later  it  was  found  that  the  metal  may 
be  prepared  by  reducing  potassium  carbonate  with  carbon 
and  by  the  electrolysis  of  potassium  chloride. 

Potassium  is  a  silver  white  metal  which  tarnishes  in- 
stantly in  moist  air  and  gives  so  much  heat  as  to  ignite  the 
liberated  hydrogen,  when  it  is  thrown  on  water.  The  hydro- 


SALTPETER  299 

gen  burns  with  a  violet  flame  characteristic  of  potassium 
and  its  compounds  when  volatilized  in  a  flame.  The  metal 
is  kept  under  kerosene  to  protect  it  from  the  action  of  the 
air. 

Potassium  Oxide,  K2O,  and  Potassium  Hydroxide,  KOH, 
are  so  similar  to  sodium  oxide  and  hydroxide  in  the  methods 
of  preparation  and  properties  that  they  require  no  separate 
description  here. 

Potassium  Chlorate,  KC1O3,  may  be  prepared  by  passing 
chlorine  into  a  warm  concentrated  solution  of  potassium 
hydroxide.  Potassium  hypochlorite,  KC1O,  is  formed  at 
first  and  this  changes  to  chlorate  and  chloride  by  self- 
oxidation.  Potassium  chlorate  is  used  in  the  laboratory 
for  the  preparation  of  oxygen.  It  is  also  used  in  medicine, 
in  flash-light  powders,  in  matches  and  in  some  of  the  primers 
used  in  shells  for  firearms. 

Potassium  Nitrate  or  Saltpeter,  KNO3. — From  the 
introduction  of  gunpowder  into  Europe,  about  1300,  it 
was  practically  the  only  explosive  used  in  hunting  and  war- 
fare till  near  the  close  of  the  nineteenth  century.  The 
saltpeter  for  the  manufacture  of  gunpowder  was  almost 
entirely  obtained,  until  comparatively  recent  times,  from 
the  decay  of  organic  matter  containing  nitrogen  and  potas- 
sium. Considerable  quantities  of  saltpeter  formed  in  this 
way  have  been  obtained  from  India.  During  the  war  of 
1812  the  United  States  depended  largely  on  saltpeter  from 
the  Mammoth  Cave,  Kentucky.  During  more  recent  times 
the  world's  supply  of  potassium  nitrate  has  been  largely 
manufactured  by  the  interaction  of  Chile  saltpeter,  NaNO3, 
from  South  America  and  potassium  chloride  from  Germany. 
With  the  introduction  of  smokeless  powder  and  other  ex- 
plosives the  demand  for  potassium  nitrate  has,  of  course, 
greatly  decreased  and  the  salt  is  no  longer  an  important 
factor  in  warfare  when  other  nitrates  can  be  obtained. 
Sodium  nitrate  is  hygroscopic  and  it  never  successfully 


300  ALKALI  METALS 

replaced  saltpeter  in  the  old  forms  of  gunpowder.  Apart 
from  that  property  it  would  be  better  than  potassium 
nitrate.  Why? 

Saltpeter  is  used  in  the  curing  of  meats,  to  which  it 
imparts  a  desirable  red  color.  Taken  in  considerable 
quantities  it  is  a  poison,  but  in  small  quantities  it  is  a  nor- 
mal constituent  of  a  number  of  vegetables. 

Gunpowder  is  a  mixture  of  about  75  parts  of  saltpeter, 
13  parts  of  charcoal  and  12  parts  of  sulfur.  This  corre- 
sponds very  nearly  to  the  equation : 

2KNO3  -f  3C  +  S  -  K2S  +  N2  +  3C02 

The  explosion  depends  on  the  rapid  formation  of  a  large 
volume  of  gases  heated  to  a  high  temperature  by  the  burning 
of  the  powder.  The  rate  of  burning  is  regulated  by  the 
size  of  the  grains,  as  these  burn  from  the  surface  inward. 
For  small  arms  a  small  size  of  grain  which  burns  very 
rapidly  is  used.  For  cannon  the  grains  must  be  large, 
sometimes  an  inch  in  diameter,  in  order  to  give  time  for 
the  heavy  ball  to  get  started  before  the  full  force  of  the 
explosive  is  developed. 

Potassium  Carbonate,  K2CO3. — The  preparation  of 
crude  potassium  carbonate  from  wood  ashes  has  been 
described.  It  is  a  deliquescent  salt,  differing  in  .this  pro- 
perty from  sodium  carbonate. 

Potassium  Bicarbonate  or  Saleratus,  KHCOs. — When 
potassium  compounds  were  more  easily  obtained  than  those 
of  sodium  this  salt  was  used  for  cooking,  but  it  has  now  been 
entirely  replaced  by  the  cheaper  sodium  salt. 

Ammonium,  NH4,  is  not,  strictly  speaking,  a  metal, 
still  less  an  element,  but  the  ammonium  salts  are  so  closely 
similar  to  the  salts  of  sodium  and  potassium  that  ammonium 
is  most  conveniently  classed  as  an  alkali  metal.  It  has 
not  been  prepared  in  the  free  state  but  when  a  strong  solu- 
tion of  ammonium  chloride  is  poured  over  sodium  amalgam 


AMMONIUM  301 

an  unstable  ammonium  amalgam  is  formed.  This  decom- 
poses quickly,  however,  into  hydrogen,  ammonia,  NH3, 
and  mercury: 

NH4Cl  +  Na(Hg)  =  NaCl  +  NH4(Hg) 

Ammonium  Hydroxide,  NH4OH. — When  ammonia  dis- 
solves in  water  two  things  occur:  the  ammonia  combines 
with  the  water  to  form  ammonium  hydroxide  and  am- 
monium hydroxide  ionizes  partly  to  ammonium,  NH4+y 
and  hydroxide,  OH~,  ions: 

NH3  +  HOH  <=>  NH4OH  +±  NH4+  +  OH~ 

Both  reactions  are  reversible  and  may  be  carried  to 
practical  completion  in  either  direction.  Ammonia  has  a 
much  lower  boiling  point  than  water,  and  ammonium 
hydroxide  dissociates  so  easily  that  from  concentrated 
solutions  nearly  pure  ammonia  escapes  on  boiling  the  solu- 
tion. This  is,  indeed,  the  most  convenient  method  of 
preparing  the  gas.  Ammonia  also  escapes  from  very  dilute 
solutions  and  all  of  the  ammonia  will  be  found  in  the  water 
which  distils  first  from  such  solutions.  This  property  is 
used  in  determining  minute  quantities  of  ammonia  in  the 
analysis  of  drinking  water.  Ammonia  derived  from  the 
organic  matter  of  sewage  may  sometimes  be  detected  in 
this  way. 

Ammonium  hydroxide  will  neutralize  strong  acids,  giving 
neutral  salts,  and  in  many  of  its  properties  it  closely  re- 
sembles the  hydroxides  of  the  alkali  metals,  though  it  is  a 
much  weaker  base.  By  this  is  meant  that  in  correspond- 
ing dilutions  there  are  many  more  hydroxide  ions  in  the 
solution  of  sodium  hydroxide  than  in  that  of  ammonium 
hydroxide. 

Ammonium  chloride,  NH^Clj  may  be  prepared  either  by 
bringing  ammonia  gas  and  hydrochloric  acid  gas  together 


302  ALKALI  METALS 

or  by  neutralizing  a  solution  of  ammonium  hydroxide  with 
dilute  hydrochloric  acid  and  evaporating  the  solution  to 
dryness : 

NH3  +  HC1  =  NH4C1 
NH4OH  +  HC1  =  NH4C1  +  HOH 

Ammonium  chloride  has  a  sharp,  salty  taste.  When 
the  dry  salt  is  heated  it  sublimes  but  at  the  same  time  dis- 
sociates into  ammonia  and  hydrochloric  acid.  This  has 
been  proved  by  the  weight  of  the  gas  which  is  formed. 

Ammonium  Salts. — A  large  number  of  ammonium  salts 
have  been  prepared.  Among  these  may  be  mentioned 
ammonium  sulfide,  (NH4)2S,  ammonium  hydrosulfide, 
NH4HS,  ammonium  sulfate,  (NH4)2S04,  ammonium  nitrate, 
NH4N03,  ammonium  carbonate,  (NH4)2C03,  ammonium 
bicarbonate,  NH4HC03,  and  ammonium  chloroplatinate, 
(NH4)2PtCl6. 

Colored  Flames.  Spectrum  Analysis. — If  a  wire  which 
has  been  dipped  in  a  strong  solution  of  salt  is  held  in  the 
blue  flame  of  a  Bunsen  burner  a  brilliant  yellow  flame  is 
produced.  A  crystal  or  solution  of  potassium  dichromate 
which  is  held  near  such  a  flame  will  be  yellow  and  the 
hand  or  face  near  such  a  flame  will  assume  a  peculiar, 
sallow  appearance  because  the  flame  illuminates  objects 
with  a  pure  yellow  light  and  none  of  the  other  colors  of 
natural  objects  can  be  brought  out  by  illumination  of  this 
kind. 

If  the  flame  is  brought  near  the  slit  of  a  spectroscope 
(Fig.  49)  at  B,  on  looking  through  the  telescope  D  a  single 
yellow  line  will  be  seen.  The  exact  position  of  the  line 
can  be  fixed  by  the  scale  E  reflected  from  the  face  of  the 
prism  A,  used  to  refract  the  light  and  separate  it  into  its 
different  colors. 

If  a  clean  wire  is  dipped  in  a  solution  of  pure  potassium 
chloride  and  then  held  in  the  flame,  a  wholly  different,  violet 


SPECTRA 


303 


flame  will  be  given.  With  the  spectroscope  this  will  give  a 
red  and  a  violet  line.  Still  different  colors  and  lines  will  be 
given  by  salts  of  calcium,  barium  and  other  metals. 

By  passing  electric  sparks  between  points  of  other  metals, 
which  cannot  be  volatilized  in  the  Bunsen  flame,  other 
characteristic  spectra  can  be  obtained  and  all  of  these  spectra 
may  be  used  to  identify  the  metals  or  elements  that  give 
them. 


FIG.  49. 


The  Solar  Spectrum. — If  the  sun  or  diffused  daylight 
is  examined  with  a  spectroscope,  a  spectrum  with  all  the 
colors  of  the  rainbow  is  seen  but  this  is  crossed  by  many 
fine  dark  lines.  These  lines  were  observed  by  Fraunhofer 
early  in  the  nineteenth  century  and  are  still  called  by  his 
name,  though  he  gave  no  explanation  for  them.  About 
1860  two  Germans,  Bunsen  and  Kirchoff,  took  up  the  study 
of  these  lines  and  showed  that  the  position  of  the  lines 
corresponds  exactly  with  the  position  of  the  bright  lines 


304  ALKALI  METALS 

in  the  spectra  of  the  elements.  Thus  there  is  a  dark  line 
in  the  solar  spectrum  which  corresponds  exactly  with  the 
yellow  line  of  the  sodium  spectrum. 

If  we  consider  that  light  is  a  vibrating  motion  sent  out 
through  space  by  a  luminous  object,  it  follows  that  there 
is  some  rapid,  periodic  motion  of  the  sodium  atoms,  or 
more  likely  of  the  electrons  within  the  sodium  atoms,  which 
sends  out  the  waves  of  yellow  light.  If  these  waves  can 
be  sent  out  by  sodium  atoms,  we  may  suppose  that  waves 
of  exactly  the  same  length  which  pass  sodium  atoms  in 
a  flame  will  be  absorbed  by  the  atoms.  That  this  is  so  can 
be  demonstrated  by  placing  a  brilliant  light  which  gives 
a  continuous  spectrum  behind  a  sodium  flame.  The  con- 
tinuous spectrum  will  now  be  crossed  by  a  dark  line  where 
the  bright  sodium  line  was  before.  The  sodium  of  the 
flame  absorbs  the  light  of  its  own  kind  from  the  more 
brilliant  light  behind  and  the  small  amount  of  sodium  light 
sent  on  appears  dark  by  contrast  with  the  more  brilliant 
light  on  both  sides  of  it  in  the  spectrum.  Something  of 
the  same  kind  doubtless  occurs  in  the  sun.  The  central 
portions  of  the  sun,  which  are  at  a  very  high  temperature, 
send  out  light  of  all  wave  lengths  and  would  give  a  continuous 
spectrum  of  all  the  primary  colors,  if  none  of  the  light  were 
absorbed.  The  atmosphere  of  the  sun,  which  is  also  at  a 
high  temperature,  contains  the  vapors  of  many  of  the 
elements  and  these  absorb  from  the  original  Iigh1>  the  vibra- 
tions which  they  themselves  emit,  with  the  result  that  less 
light  of  those  wave  lengths  corresponding  to  the  vibrations 
of  the  atoms  goes  through.  This  gives  dark  lines  in  the 
solar  spectrum  corresponding  to  the  bright  lines  of  sodium, 
iron  and,  in  all,  of  about  thirty  elements  which  are  found 
on  the  earth.  The  presence  of  hydrogen,  helium  and 
other  elements  has  also  been  demonstrated  in  many  of  the 
stars  so  far  distant  that  light  requires  many  years  for  its 
passage  from  them  to  the  earth. 


SUMMARY.    ALKALI  METALS  305 

SUMMARY 

The  alkali  metals  stand  at  the  extreme  of  metallic  prop- 
erties, as  the  halogens  are  at  the  extreme  of  the  non-metallic 
elements.  They  decompose  water  at  ordinary  temperatures 
and  their  salts  are  easily  soluble  in  aqueous  solutions. 

Sodium  occurs  chiefly  as  common  salt  but  also  in  most 
natural  silicates  and  in  Chile  saltpeter. 

Sodium  chloride  is  very  widely  diffused  in  nature  and  is  a 
necessary  constituent  of  foods.  It  is  the  source  of  nearly 
all  compounds  of  sodium  used  in  the  industries. 

Chile  saltpeter  has  been  the  chief  source  of  nitric  acid 
for  the  manufacture  of  explosives  and  for  other  uses. 

Sodium  carbonate  has  been  made  by  the  Le  Blanc  and 
by  the  ammonia-soda  process.  The  preparation  of  sodium 
hydroxide  by  the  electrolysis  of  salt  is  taking  the  place  of 
these  processes  in  considerable  measure. 

The  Le  Blanc  process  consists  of  three  steps:  the  prepa- 
ration of  sodium  sulfate  from  salt,  reduction  to  sodium 
sulfide  and  conversion  to  sodium  carbonate  carried  out 
simultaneously,  and  separation  of  calcium  sulfide  from  the 
sodium  carbonate. 

Sodium  carbonate  is  hydrolyzed  by  water  because  car- 
bonic acid  is  a  weak  acid  and  the  bicarbonate  ion,  HCOa", 
ionizes  scarcely  more  than  ^water  itself. 

Sodium  hydroxide  is  prepared  by  treating  sodium  car- 
bonate with  milk  of  lime  or  by  the  electrolysis  of  a  solution 
of  salt. 

Sodium  peroxide  is  prepared  by  heating  sodium  in  air. 
It  is  used  in  helmets  for  mine  rescue  work  and  for  the  prepa- 
ration of  hydrogen  peroxide  to  use  in  bleaching  silks  and 
wool. 

Sodium  sulfite  is  a  reducing  agent,  used  in  photographic 
developers. 

Soluble  glass  is  used  as  a  fi reproofing  agent  and  to  preserve 
eggs. 

20 


306  ALKALI  METALS 

Borax  is  -used  in  washing  and  for  welding.. 

Potassium  is  found  in  many  natural  silicates,  in  wood 
ashes  and  in  mineral  deposits  or  solutions  in  Germany, 
France,  California  and  elsewhere.  It  is-  a  necessary  con- 
stituent of  fertile  soils. 

Soft  soap  was  formerly  made  with  the  lye  from  leaching 
wood  ashes. 

Metallic  potassium  was  first  prepared  by  Sir  Humphrey 
Davy  by  the  electrolysis  of  potassium  hydroxide. 

Potassium  oxide  and  hydroxide  are  prepared  in  the 
same  manner  as  the  corresponding  sodium  compounds. 

Potassium  chlorate  is  used  in  medicine,  in  matches,  primers 
and  flash-light  powders. 

Potassium  nitrate  is  used  in  the  manufacture  of  black 
gunpowder.  Smokeless  powders  have  displaced  this  use 
for  warfare. 

Gunpowder  is  a  mixture  of  charcoal,  sulfur  and  saltpeter. 

Potassium  carbonate  was  formerly  obtained  from  wood 
ashes  but  is  now  manufactured  from  other  potassium 
salts. 

Potassium  bicarbonate  was  formerly  used  in  cooking, 
under  the  name  of  saleratus,  but  has  been  replaced  by  "  cook- 
ing soda." 

Ammonium  is  only  known  in  salts  and  as  a  very  dilute 
amalgam. 

Ammonium  hydroxide  is  a  weak  base,  partly  because  of 
slight  ionization  and  partly  because  of  dissociation  to 
ammonia  and  water. 

Ammonium  chloride  is  prepared  by  the  direct  union  of 
ammonia  and  hydrochloric  acid  or  by  neutralizing  ammo- 
nium hydroxide  with  hydrochloric  acid.  Many  other  am- 
monium salts  are  known. 

Some  elements  impart  characteristic  colors  to  the  flame  of 
a  Bunsen  burner  and  many  others  develop  luminous  flames 
with  the  electric  spark.  By  means  of  the  spectroscope  such 


.EXERCISES.     ALKALI  METALS  307 

elements  may  'be  identified,  even  when  they  are  present  in 
very  minute  quantities. 

Vapors  of  elements  absorb  and  diffuse  light  of  the  same 
wave  lengths  as  those  of  the  light  which  they  emit.  In  this 
manner  the  vapors  in  the  corona  of  the  sun  cause  the  dark 
lines  of  the  solar  spectrum. 

EXERCISES 

1.  Write  the  equations  for  the  reactions  between  sodium  oxide 
and  sodium  peroxide  and  hydrochloric  acid. 

2.  In  the  presence  of  a  little  copper  oxide  as  a  catalytic  agent 
sodium  peroxide  dissolves  in  water  with  the  evolution  of  oxygen. 
Write  the  equation.     The  fused  sodium  peroxide  containing  cop- 
per oxide  is  called  "oxone. " 

3.  Sodium    thiosulfate   decomposes   with   acids,  giving  sulfur 
dioxide  and  free  sulfur.     Write  the  equation* 

4.  Write  the  equations  for  the  domestic  manufacture  of  soft 
soap,  assuming  that  the  -grease  used  consists  chiefly  of  stearin. 

5.  Write  the  equation  for  the  preparation  of  potassium  chlorate. 

6.  Write   the   equation  for  the  decomposition  of  ammonium 
nitrate  by  heat. 

7.  What  is  the  weight  of  a  gram-molecular  volume  of!  the  gases 
obtained  by  heating  ammonium  chloride? 


CHAPTER  XXVII 

GROUP  I:  SECOND  DIVISION;  COPPER,  SILVER,  GOLD; 
PHOTOGRAPHY 

Contrast  between  the  First  and  Second  Divisions  of 
Group  I. — There  is  a  very  marked  contrast  between  sodium 
and  potassium  on  the  one  hand  and  copper,  silver  and  gold 
on  the  other;  so  great,  indeed,  that  the  metals  are  classified 
in  the  same  group  only  on  account  of  their  valences  and 
their  positions  in  the  periodic  system. 

Potassium  and  sodium  are  very  light  metals.  They  are 
extremely  active,  tarnishing  instantly  in  moist  air,  and  de- 
composing water  energetically  at  ordinary  temperatures. 
Their  hydroxides  are  easily  soluble  and  are  very  strong 
bases. 

Copper,  silver  and  gold  are  heavy  metals,  having  a 
brilliant  metallic  luster,  and  only  copper  is  affected  by 
moist  air.  They  do  not  decompose  water  at  any  tempera- 
ture and  are  scarcely  affected  even  by  cold  dilute  hydro- 
chloric or  sulfuric  acid. 

The  closest  resemblance  to  the  first  division  of  the  group 
is  found  in  the  chlorides,  Cu2Cl2  (CuCl),  AgCl  and  AuCl 
and  in  the  series  of  oxides,  Cu2O,  Ag2O  and  Au2O,  but  the 
oxides  CuO,  Ag2O2  and  Au2O3  are  also  known. 

Copper.  Occurrence,  Metallurgy. — The  most  important 
minerals  containing  copper  are  chalcopyrite,  or  copper 
pyrites,  CuFeS2,  chalcocite,  Cu2S,  and  malachite,  CuCO3.- 
Cu(OH)2.  It  is  also  found  as  the  free  metal,  especially 
in  the  Lake  Superior  region.  The  sulfide  ores  are  most 
common  and  from  these  the  copper  is  first  separated  as  a 

308 


COPPER  809 

mixture  of  cuprous  sulfide,  Cu2S,  and  ferrous  sulfide,  FeS, 
which  is  called  copper  matte  and  which  melts  at  a  low  tem- 
perature. The  copper  matte  is  then  roasted  in  an  appa- 
ratus similar  to  the  Bessemer  converter  (p.  327),  the  sulfur 
escaping  as  sulfur  dioxide.  The  ferrous  oxide  combines 
with  fine  sand  or  silica,  which  is  added,  to  form  ferrous 
silicate  and  the  copper  is  reduced  to  the  metallic  state. 

The  copper  obtained  by  the  process  which  has  been  out- 
lined is  quite  impure,  containing,  usually,  some  gold  and 
silver  and  larger  amounts  of  arsenic,  lead  and  other  metals. 
The  arsenic,  especially,  greatly  reduces  the  conductivity  of 
copper  for  electricity  and  renders  it  unfit  for  many  industrial 
uses.  This  crude  copper  is  refined  electrolytically  by  sus- 
pending plates  of  it  in  a  solution  of  copper  sulfate  and  pass- 
ing an  electric  current  from  the  plates  to  a  cathode  of  pure 
copper.  The  anode  of  crude  copper  dissolves  and  nearly 
pure  copper  is  deposited  on  the  cathode. 

Copper  is  a  red  metal,  which  does  not  tarnish  in  dry 
air.  It  is  blackened  by  hydrogen  sulfide  and  when  exposed 
to  the  weather  it  becomes  covered  with  a  coating  of 
basic  carbonate,  which  has  the  composition  of  malachite, 
CuCO3.Cu(OH)2.  Copper  is  scarcely  attacked  by  cold, 
dilute  hydrochloric  or  sulfuric  acid  but  dissolves  easily  in 
nitric  acid. 

Copper  is  used  for  electrical  conductors,  for  the  sheathing 
of  ships  and  the  manufacture  of  brass,  bronze  and  other 
alloys. 

Salts  of  Copper. — Copper  forms  cuprous  compounds,  such 
as  Cu2O,  Cu2S  and  Cu2Cl2,  in  which  it  appears  to  be  uni- 
valent,  and  cupric  compounds,  such  as  CuO,  CuS,  CuCl2 
and  CuS04,  in  which  it  is  bivalent. 

Cuprous  Oxide,  Cu2O. — If  sodium  hydroxide  is  added 
to  a  solution  of  copper  sulfate  containing  Rochelle  salt 
(sodium  potassium  tartrate),  no  precipitate  of  copper 
hydroxide  is  formed  because  the  copper  forms  a  complex  ion 


310          COPPER,  SILVER,  GOLD;  PHOTOGRAPHY 

with  the  tartrate  and  there  are  very  few  copper  ions,  Cu++, 
in  the  solution.  On  warming  the  solution  and  adding  a 
little  glucose  the  copper  is  reduced  to  the  cuprous  form  and 
separates  as  red  cuprous  oxide.  This  process  is  used  as 
a  method  for  the  detection  and  determination  of  glucose 
(p.  223)  in  sugar  analysis  and  for  diagnosis. 

Cuprous  Chloride,  Cu2Cl2,  is  a  white  compound  almost 
insoluble  in  water  but  more  easily  soluble  in  concentrated 
hydrochloric  acid.  It  is  easily  prepared  by  digesting  copper 
turnings  with  cupric  chloride,  CuCl2,  and  concentrated 
hydrochloric  acid.  The  cuprous  chloride  is  precipitated  on 
adding  water  to  the  solution. 

Copper  Sulfate  or  Blue  Vitriol,  CuSO4.5H2O,  is  the  most 
common  and  best  known  of  the  salts  of  copper.  It  may  be 
prepared  by  dissolving  copper  oxide  in  dilute  sulfuric  acid 
or  by  dissolving  metallic  copper  in  hot  concentrated  sulfuric 
acid.  It  is  used  in  the  electrolytic  refining  of  copper,  in 
electrotyping  and  electroplating,  as  a  mordant  in  dyeing, 
and  in  the  gravity  cells  which  were  formerly  much  used  for 
telegraphic  purposes.  A  mixture  of  copper  sulfate,  slaked 
lime  and  water,  called  Bordeaux  mixture,  is  used  for  spray- 
ing fruit  trees. 

Silver.  Occurrence,  Metallurgy. — Silver  is  sometimes 
found  in  the  free  state  in  nature  but  more  often  it  occurs  as 
the  sulfide,  usually  associated  with  other  sulfides,  especially 
with  galena,  or  lead  sulfide.  When  galena  is  reduced  to 
metallic  lead  the  silver  is  also  reduced.  From  the  lead  con- 
taining silver  the  metal  is  recovered  by  melting  the  lead  and 
mixing  it  thoroughly  with  a  comparatively  small  amount 
of  zinc.  On  allowing  the  melted  mixture  to  stand,  nearly 
all  of  the  zinc  rises  to  the  top,  carrying  the  silver  with  it. 
The  zinc  is  then  skimmed  off  and  the  alloy  of  silver  and  zinc 
with  a  little  lead  is  heated  in  a  retort  to  distil  away  the  zinc. 
The  silver  and  lead  which  remain  are  then  heated  in  the  air 
to  oxidize  the  lead,  leaving  almost  pure  silver  behind. 


SILVER;  PHOTOGRAPHY  311 

Silver  is  a  white  metal  which  does  not  tarnish  in  dry  or 
moist  air  at  any  temperature.  It  is  easily  blackened  by 
hydrogen  sulfide.  It  dissolves  in  dilute  nitric  acid  or  in  hot 
concentrated  sulfuric  acid,  very  much  as  copper  does,  but 
it  is  univalent  in  the  salts  formed,  while  copper  is  bivalent. 

Silver  coins  of  the  United  States  contain  90  per  cent  of 
silver  and  10  per  cent  of  copper,  the  copper  being  added  to 
give  hardness  and  to  lessen  the  wear  of  use. 

Silver  Nitrate  or  Lunar  Caustic,  AgNO3,  is  prepared  by 
dissolving  silver  in  nitric  acid.  The  salt  melts  easily  and 
sticks  of  the  fused  salt  are  used  under  the  name  of  lunar 
caustic  to  cauterize  wounds.  The  name  dates  from  the 
time  of  the  alchemists  when  silver  was  associated  in  their 
literature  with  the  moon  and  gold  with  the  sun. 

Silver  Chloride,  AgCl,  is  formed  as  a  curdy,  white  pre- 
cipitate when  a  solution  of  salt  or  of  hydrochloric  acid 
is  added  to  a  solution  of  silver  nitrate.  Silver  bromide, 
AgBr,  is  •  a  yellowish-while  salt  and  silver  iodide,  Agl, 
is  a  yellow  salt.  Each  of  these  may  be  prepared  in  the 
same  manner  as  silver  chloride.  Each  of  these  salts  is 
sensitive  to  the  action  of  light,  being  reduced  to  metallic 
silver  or  to  a  salt  containing  less  of  the  halogen. 

Silver  chloride  dissolves  easily  in  ammonia  and  each  of  the 
halogen  salts  dissolves  in  a  solution  of  sodium  thiosulfate 
("hyposulfite"),  Na2S2O3. 

Photography. — The  sensitive  character  of  the  silver  halides 
when  exposed  to  the  light  is  used  in  photography.  "Dry 
plates"  are  thin  sheets  of  celluloid  or  plates  of  glass  which 
have  been  covered  with  a  thin  coating  of  an  emulsion  of 
silver  bromide,  AgBr,  in  a  solution  of  gelatin.  The  plate  is 
exposed  in  a  camera  for  a  moment  to  the  image  of  the  object 
which  is  to  be  photographed.  After  this  exposure  the  plate 
appears  to  the  eye  entirely  unchanged,  but  in  some  manner, 
at  present  only  very  vaguely  understood,  some  of  the  mole- 
cules of  the  silver  bromide  are  changed  by  the  light  and 


312          COPPER,  SILVER,  GOLD;  PHOTOGRAPHY 

are  sensitive  to  the  action  of  a  reducing  agent  called  a 
"developer."  After  "developing"  the  picture,  or  simul- 
taneously with  the  development,  the  portions  of  the  silver 
bromide  which  have  not  been  affected  are  dissolved  and 
removed  by  a  solution  of  sodium  thiosulfate  ("hypo- 
sulfite"  or  "hypo").  This  last  process  is  called  fixing. 
Without  this  treatment  the  whole  plate  would  become 
black. 

In  the  picture  obtained  in  this  way  the  lights  and  shadows 
are,  of  course,  reversed.  The  portions  of  the  object  which 
were  light  appear  dark  in  the  "negative."  A  "positive" 
picture  is  obtained  by  placing  the  negative  over  a  sheet  of 
sensitive  paper  and  exposing  it  to  the  light.  The  picture 
printed  in  this  manner  is  fixed,  as  before,  by  soaking  the 
paper  in  a  solution  of  sodium  thiosulfate  to  remove  the 
unchanged  silver  salt. 

Gold.  Occurrence,  Metallurgy. — Gold  is  found  almost 
exclusively  in  the  free  state.  Occasionally  large  nuggets 
weighing  many  pounds  and  worth  thousands  of  dollars  have 
been  found,  but  such  nuggets  are  extremely  rare.  Usually 
gold  is  found  in  small  grains  mixed  with  sand  or  gravel  or 
disseminated  in  quartz,  pyrite  and  other  rocks  and  minerals. 
There  is  a  very  minute  quantity  of  gold,  worth,  perhaps,  two 
or  three  cents,  in  a  ton  of  sea  water,  but  no  profitable  method 
of  recovering  it  has  been  discovered. 

Gold  has  a  specific  gravity  of  19.26  and  because  it  is  so 
heavy  the  comparatively  light  minerals  composing  sand 
or  gravel  can  be  separated  from  it  easily  in  a  current  of 
water.  For  rich  sands  and  gravels  the  process  may  be 
carried  out  by  hand,  with  a  pan,  from  which  comes  the 
expression  "to  pan  out."  It  is  also  carried  out  on  a  large 
scale  in  "hydraulic  mining." 

For  massive  rocks  the  cyanide  process  is  now  most  com- 
mon. The  ore  is  first  broken  to  a  powder  in  the  stamping 
mills  and  the  gold  and  silver  are  extracted  by  a  solution  of 


GOLD  313 

potassium  or  sodium  cyanide.     Oxygen  from  the  air  or  some 
oxidizing  agent  is  required  in  the  solution  of  the  gold : 

4Au  +  8KCN  +  O2  +  2H2O  =  4KAu(CN)2  +  4KOH 

From  the  double  cyanide  of  potassium  and  gold  the  latter 
is  easily  precipitated  by  metallic  zinc. 

Native  gold  and  the  gold  obtained  by  the  cyanide  or 
other  processes  practically  always  contains  silver.  To 
separate  the  two  metals  enough  silver  is  added,  if  necessary, 
to  give  an  alloy  containing  not  more  than  one-third  of 
its  weight  in  gold.  After  melting  the  mixture  the  alloy 
is  treated  with  nitric  acid  or  with  hot,  concentrated  sulfuric 
acid,  either  of  which  dissolves  the  silver,  leaving  the  gold 
nearly  pure. 

English  gold  coins  are  22  carats  fine,  i.e.,  2%4  pure  gold. 
The  United  States  coin  is  900  fine,  containing  900  parts  of 
gold  to  100  parts  of  copper. 

Gold  does  not  dissolve  in  either  of  the  three  common 
acids,  hydrochloric,  nitric  or  sulfuric,  but  it  dissolves  readily 
in  aqua  regia. 

Gold  forms  three  oxides,  Au2O,  AuO  and  Au203;  and 
three  chlorides,  AuCl,  AuCl2  and  AuCl3. 

SUMMARY 

The  second  division  of  Group  I  contains  copper,  silver 
and  gold.  These  do  not  decompose  water  at  any  tempera- 
ture and  differ  very  greatly  from  sodium  and  potassium. 
They  are  univalent  in  only  a  part  of  their  compounds. 

Copper  is  found  as  chalcopyrite,  chalcocite  and  malachite. 
The  sulfide  ores  are  concentrated  to  copper  matte  and  the 
latter  is  reduced  by  heating  it  with  sand  in  a  blast  of  air. 

Copper  dissolves  in  nitric  acid  or  in  hot,  concentrated 
sulfuric  acid,  but  it  is  not  affected  by  hydrochloric  acid  or 
dilute  sulfuric  acid  in  the  absence  of  air. 


314          COPPER,  SILVER,  GOLD;  PHOTOGRAPHY 

Copper  forms  two  oxides,  two  chlorides  and  two  sulfides. 
The  sulfate,  blue  vitriol,  is  the  most  common  salt. 

Red  cuprous  oxide  is  formed  by  the  reduction  of  a  solution 
of  a  copper  salt  in  an  alkaline  tart  rate  by  means  of  glucose. 

Cuprous  chloride  is  prepared  by  reducing  cupric  chloride 
by  means  of  copper  in  the  presence  of  concentrated  hydro- 
chloric acid.  It  is  almost  insoluble  in  water  or  dilute  acids. 

Blue  vitriol  is  prepared  by  dissolving  copper  oxide  in 
dilute  sulfuric  acid. 

Silver  is  separated  from  lead  by  adding  zinc  to  the  melted 
metal,  nearly  all  of  the  silver  dissolving  in  the  zinc  and  rising 
to  the  top. 

Silver  resembles  copper  in  its  conduct  toward  nitric  or 
sulfuric  acid. 

Silver  coins  are  an  alloy  of  silver  and  copper. 

Silver  nitrate  or  lunar  caustic  is  prepared  by  dissolving 
silver  in  nitric  acid. 

Silver  chloride,  silver  bromide  and  silver  iodide  are 
insoluble. 

In  photographic  dry  plates  silver  bromide  is  so  affected  by 
light  that  it  is  easily  reduced  to  metallic  silver.  The  un- 
changed silver  bromide  is  removed  by  a  solution  of  sodium 
thiosulfate. 

Gold  occurs  in  the  free  state  and  is  obtained  by  washing 
away  lighter  minerals  with  water  or  by  dissolving  it  in  a 
solution  of  potassium  cyanide  and  precipitating  it  with  zinc. 

Gold  coins  are  alloys  of  gold  and  copper. 

Gold  does  not  dissolve  in  any  of  the  common  acids.  It 
dissolves  in  aqua  regia. 

Gold  forms  three  oxides  and  three  chlorides. 

EXERCISES 

1.  Write  the  equation  for  the  reduction  of  copper  matte  (p.  309) . 
Ferrous  silicate  is  Fe2Si04. 

2.  Write  the  equations  for  the  solution  of  copper  and  of  silver 
in  nitric  acid.     Nitric  oxide  is  evolved. 


EXERCISES.     COPPER,  SILVER,  GOLD  315 

3.  Write  the  equations  for  the  solution  of  copper  and  of  silver 
in  concentrated  sulfuric  acid.     Sulfur  dioxide  is  formed. 

4.  How  much  silver  and  how  much  concentrated  sulfuric  acid 
would  be  required  to  give  22.4  liters  of  sulfur  dioxide?     One  liter 
of  sulfur  dioxide  weighs  2.93  grams. 

6.  How  much  chloroauric  acid,  HAuCl-j,  can  be  prepared  from 
10  grams  of  gold? 


CHAPTER  XXVIII 

GROUP  VI.     SECOND  DIVISION:    CHROMIUM,  TUNG- 
STEN, URANIUM 

Classification  of  Chromium,  Tungsten  and  Uranium. — 
The  three  metals  considered  in  this  chapter  belong  to  the 
second  division  of  Group  VI.  The  compounds  potassium 
chromate,  K2C  O4,  and  sodium  tungstate,  Na2WO4,  show 
a  close  relationship  to  potassium  sulfate,  K2SO4,  but  while 
there  is  this  relationship  in  some  of  their  compounds  the 
free  elements  show  very  strong  contrasts.  Sulfur  is  a  solid 
which  melts  at  a  temperature  only  a  little  above  the  boiling 
point  of  water.  It  is  a  typical  non-metallic  element,  both 
in  the  free  state  and  in  its  compounds.  Chromium,  tung- 
sten and  uranium,  on  the  other  hand,  are  metals  with  very 
high  melting  points. 

Chromium.  Occurrence,  Metallurgy. — Chromium  is 
found  chiefly  in  nature  as  the  mineral  chromite  or  chrome 
iron  ore,  FeCr2O4.  The  mineral  is  isomorphous  with 
magnetite,  Fe->O4.  In  other  words,  the  two  minerals  have 
the  same  crystalline  form,  indicating  a  similar  chemical 
structure. 

Metallic  chromium  is  prepared  by  igniting  a  mixture  of 
chromic  oxide  and  aluminium : 

Cr203  +  2A1  =  2Cr  +  A1203 

Pure  chromium  is  a  very  hard,  brittle,  crystalline  metal. 
It  is  used  as  an  addition  to  steel  to  give  it  extreme  hardness 
and  toughness  for  use  in  the  armor  plate  of  battleships. 

316 


CHROMIUM  317 

Potassium  Chromate,  K2CrO4,  and  Potassium  Bichro- 
mate, K2Cr2O7. — When  chrome  iron  ore  is  heated  in  the  air 
with  potassium  carbonate  or  potassium  hydroxide  it  is 
slowly  decomposed  and  oxidized,  the  iron  giving  ferric  oxide, 
Fe203,  and  the  chromium  forming  potassium  chromate, 
K2CrO4.  Potassium  chromate  is  very  easily  soluble  in 
water.  The  salt  and  its  solution  have  a  clear,  lemon- 
yellow  color.  Acids  convert  it  into  potassium  dichromate, 
K2Cr2O7,  a  reddish-orange  salt  which  is  much  less  soluble 
than  potassium  chromate  and  which  crystallizes  easily. 

Potassium  dichromate  is  the  best  known  compound  of 
chromium.  It  is  used  as  the  starting  point  for  the  prepa- 
ration of  nearly  all  other  compounds  of  the  element.  It  is 
also  used  in  acid  solutions  as  an  oxidizing  agent.  It  is  an 
important  mordant  for  dyeing,  and  is  employed  in  chrome 
tanning. 

Chromic  Anhydride,  CrO3. — When  concentrated  sulfuric 
acid  is  added  to  a  strong  solution  of  potassium  dichromate, 
chromic  anhydride,  which  is  not  very  soluble  in  the  con- 
centrated acid,  separates  in  red  needles.  The  compound 
corresponds  to  sulfuric  anhydride,  SO3.  It  is  a  vigorous 
oxidizing  agent. 

Lead  Chromate  or  Chrome  Yellow,  PbCrO4,  is  a  yellow, 
insoluble  salt  formed  by  adding  either  a  solution  of  potas- 
sium chromate,  K2Cr04,  or  one  of  potassium  dichromate, 
K2Cr207,  to  a  solution  of  sugar  of  lead,  Pb(C2H3O2)2.  It 
forms  an  excellent  yellow  paint. 

Chrome  Alum,  KCr(SO4)2.12H2O.— If  alcohol,  C2H6O, 
is  added  to  a  solution  of  potassium  dichromate  and  sulfuric 
acid  it  is  oxidized  to  aldehyde,  C2H4O,  while  the  chromium 
and  potassium  combine  with  the  sulfuric  acid.  Chrome 
alum,  which  is  rather  easily  soluble,  may  be  obtained  by 
evaporating  the  solution.  The  salt  is  isomorphous  with 
ordinary  alum,  indicating  a  similar  arrangement  of  the 
atoms  in  the  two  kinds  of  crystals. 


318  CHROMIUM,  TUNGSTEN,  URANIUM 

Chromic  Hydroxide,  Cr(OH)3,  and  Chromic  Oxide, 
Cr2O3. — The  addition  of  ammonium  hydroxide,  NH4OH, 
to  a  solution  of  chrome  alum  or  of  almost  any  chromic 
salt  produces  a  light  green  precipitate  of  chromic  hydrox- 
ide, Cr(OH)3.  If  the  precipitate  is  separated  and  heated, 
it  decomposes  into  water  and  chromic  oxide,  C^Os,  which 
forms  a  dark  green  powder. 

Chrome  Tanning. — The  older  methods  of  tanning  leather 
require  the  treatment  of  hides  with  the  solution  of  tannin 
extracted  from  oak  or  hemlock  bark  or  obtained  from  other 
sources.  The  process  is  slow  and  tedious,  lasting  for  several 
months.  Another  method  of  tanning  has  been  developed 
which  consists  in  treating  the  hides  first  with  potassium 
dichromate  and  dilute  sulfuric  acid,  and  then  with  acid 
sodium  sulfite,  NaHSO3,  which  causes  the  precipitation  of 
chromic  hydroxide  in  the  fiber. 

Tungsten, — The  use  of  metallic  tungsten  for  the  filaments 
of  electric  light  bulbs  has  made  the  name  of  the  element 
almost  a  household  word.  The  discovery  of  this  use  was 
made  by  a  chemist  who  was  Jed  to  it  from  noticing  the  posi- 
tion of  the  element  in  the  periodic  table. 

Tungsten  is  also  used  in  the  high-speed  tool  steels. 
Ordinary  steel  loses  its  temper  and  becomes  soft  when  it  is 
heated  and  such  steel  cannot  be  used  successfully  in  rapid 
lathe  work  when  the  cutting  tool  often  becomes  very  hot. 
By  adding  tungsten,  tools  are  made  which  retain  their  tem- 
per at  comparatively  high  temperatures.  It  is  not  too  much 
to  say  that  this  discovery  has  been  worth  many  millions  of 
dollars  in  the  saving  of  time  for  expensive  machinery  and 
skilled  workmen. 

SUMMARY 

Chromium,  tungsten  and  uranium,  of  the  second  divi- 
sion of  Group  VI,  are  metals  with  high  melting  points  and 
both  in  the  free  state  and  in  those  compounds  in  which 


SUMMARY.     CHROMIUM  319 

they  have  a  valence  of  two  they  are  distinctly  metallic, 
differing  very  markedly  from  oxygen  and  sulfur,  of  the  first 
division  of  the  same  group. 

Chromium  occurs  as  chromite.  It  is  reduced  from  its 
oxide  by  means  of  aluminium. 

Chrome  steel  is  used  for  armor  plate. 

Potassium  chromate  is  prepared  by  heating  chromite 
with  potassium  hydroxide  or  carbonate  in  contact  with  air. 
Acids  change  it  to  potassium  dichromate. 

Potassium  dichromate  is  used  with  sulfuric  acid  as  an 
oxidizing  agent.  It  is  also  used  as  a  mordant  in  dyeing. 

Lead  chromate,  or  chrome  yellow,  is  used  as  a  paint. 

Chrome  alum  is  formed  by  the  reduction  of  potassium 
dichromate  by  alcohol  in  the  presence  of  sulfuric  acid. 

Chrome  green  is  a  mixture  of  chrome  yellow  and  Prussian 
blue. 

Hides  are  tanned  by  treating  them  first  with  potassium 
dichromate  and  sulfuric  acid  and  then  with  acid  sodium 
sulfite. 

Tungsten  is  used  in  the  filaments  for  electric  light  bulbs 
and  in  the  steel  tools  for  high-speed  lathes. 

EXERCISES 

Write  equations  for  the  following: 

1.  Potassium  chromate  from  chrome  iron  ore. 

2.  Preparation  of  chrome  alum. 

3.  Preparation  of  chrome  yellow. 

4.  Chrome  tanning. 


CHAPTER  XXIX 
GROUP  VII.  SECOND  DIVISION:  MANGANESE 

Group  VII. — Manganese  is  the  only  element  belonging 
to  the  second  division  of  Group  VII  which  has  thus  far 
been  discovered,  unless  some  of  the  radioactive  elements 
with  very  brief  life  periods  belong  here.  We  should  expect 
elements  with  atomic  weights  of  about  100,  146  and  190. 
There  should  also  be  a  halogen  belonging  to  the  first  division 
of  the  group  and  having  an  atomic  weight  of  220.  It 
seems  possible  that  atoms  of  Group  VII  with  these  atomic 
weights  are  unstable  because  of  the  structure  of  such  atoms. 
According  to  the  electron  theory,  halogen  atoms  and  man- 
ganese have  seven  movable  electrons  which  they  lose  by 
transfer  to  the  oxygen  atoms  in  such  compounds  as  HC1O4, 
KMnO4  and  Mn2O7,  and  the  loss  of  many  electrons  may 
render  the  atoms  of  higher  atomic  weight  than  manganese 
and  iodine  unstable.  Such  an  explanation  is,  however, 
scarcely  more  than  an  interesting  speculation  at  the  present 
time. 

Manganese  resembles  chlorine  in  the  compounds  MnO2, 
HMn04  and  Mn2O7.  In  the  compounds  in  which  it  has  a 
valence  of  two  or  three  it  is  more  metallic  in  character  and 
is  much  more  closely  related  to  iron  than  to  the  halogens. 

Manganese.  Occurrence,  Uses. — Manganese  is  found  in 
its  purer  ores  chiefly  as  manganese  dioxide,  Mn02.  Practi- 
cally all  iron  ores  contain  some  manganese,  though  it  is 
usually  present  in  only  small  amounts.  During  the  first  half 
of  the  nineteenth  century  manganese  came  into  extensive 
use  in  the  manufacture  of  chlorine  for  bleaching  purposes. 
As  the  ore  became  scarce  and  expensive,  manufacturers 

320 


COMPOUNDS  OF  MAGNESIA  321 

of  chlorine  introduced  methods  for  the  recovery  of  the 
manganese  and  conversion  into  a  form  which  could  be 
used  over  and  over  in  the  process.  In  this  manner  the 
demand  for  manganese  dioxide  was  greatly  reduced.  During 
comparatively  recent  times  the  electrolytic  processes  for 
making  chlorine  are  rapidly  displacing  this  use  of  com- 
pounds of  manganese  altogether. 

Shortly  after  the  middle  of  the  nineteenth  century  Besse- 
mer invented  a  process  for  manufacturing  steel  which  makes 
use  of  a  form  of  cast  iron  called  spiegeleisen  (German  for 
" mirror  iron"),  because  of  its  brilliant  white  luster.  This 
form  of  iron  usually  contains  eight  or  ten  per  cent  of  man- 
ganese. Another  alloy  of  iron  and  manganese,  called  ferro- 
manganese  and  containing,  sometimes,  seventy-five  per 
cent  or  more  of  the  metal,  is  extensively  used  in  castings  for 
car  wheels  and  in  the  manufacture  of  other  products  made  of 
iron  and  steel.  These  newer  uses  have  created  a  great 
demand  for  ores  of  manganese,  and  iron  manufacturers  are 
searching  the  world  over  for  new  sources  of  supply. 

Spiegeleisen  and  ferromanganese  are  made  in  blast  fur- 
naces by  the  same  processes  that  are  used  in  making  iron 
(p.  324). 

Oxides  of  Manganese. — There  are  no  less  than  five 
different  oxides  of  manganese.  The  best  known  and  most 
important  is  manganese  dioxide,  MnO2,  sometimes  called 
black  oxide  of  manganese  because  of  its  intense  black  color. 
Its  use  in  the  manufacture  of  chlorine  has  been  referred  to 
above.  It  is  also  used  as  the  depolarizing  agent  in  dry 
batteries. 

Potassium  Manganate,  K2MnO4,  and  Potassium  Per- 
manganate, KMnO4. — If  a  mixture  of  manganese  dioxide 
and  potassium  carbonate  is  heated  in  the  air  a  green  mass 
containing  potassium  manganate  is  formed,  much  as  potas- 
sium chromate  is  formed  from  chrome  iron  ore.  The 
oxidation  may  be  hastened  by  using  potassium  nitrate  or 


322  MANGANESE 

potassium  chlorate  as  an  oxidizing  agent.  The  green  mass 
dissolves  in  water  to  a  green  solution  but  if  this  is  boiled, 
especially  if  it  is  made  slightly  acid,  a  part  of  the  manga- 
nese is  reduced  to  manganese  dioxide  while  another  part  is 
oxidized  to  potassium  permanganate,  which  gives  a  solu- 
tion with  an  intense  red  color.  It  will  be  recalled  that 
several  compounds  of  chlorine  show  a  similar  self  oxidation 
and  reduction. 

Potassium  permanganate  is  used  as  an  oxidizing  agent  in 
the  laboratory.  It  is  also  a  very  efficient  germicide  and 
disinfectant.  If  quickly  applied  to  the  wound,  it  is  the 
best  antidote  known  for  snake  bite. 

SUMMARY 

Manganese  is  the  only  known  element  of  the  second 
division  of  Group  VII.  It  resembles  chlorine  in  some  of 
its  compounds  but  it  also  resembles  iron. 

Manganese  occurs  as  the  dioxide  and  as  a  constituent  of 
nearly  all  iron  ores.  Its  alloys  with  iron  are  extensively 
used  in  the  manufacture  of  iron  and  steel. 

Manganese  dioxide  was  formerly  used  in  the  manufacture 
of  chlorine.  It  is  used  as  a  depolarizer  in  dry  cells. 

Potassium  permanganate  is  used  as  a  disinfectant  and  as 
an  oxidizing  agent. 

EXERCISES 

Write  the  equations  for  the  following: 

1.  Potassium  manganate  from  manganese  dioxide,  potassium 
carbonate  and  air. 

2.  Potassium  permanganate,  manganese  dioxide  and  potassium 
hydroxide  from  potassium  manganate  and  water. 

3.  Manganese  heptoxide  from  potassium  permanganate  and 
sulfuric  acid. 

4.  How  much  potassium  permanganate  must  be  dissolved  in  a 
liter  of  water  to  give  a  solution  of  which  1  cc.  will  yield  8  mg. 
of  oxygen  when  the  potassium  permanganate  is  reduced  to  man- 
ganese sulfate,  MnS04,  in  the  presence  of  sulfuric  acid? 


CHAPTER  XXX 
GROUP  VIII:  IRON,  COBALT,  NICKEL,  PLATINUM 

Relation  of  Group  VIII  to  Other  Elements. — In  the  midst 
of  each  of  the  longer  periods  of  the  periodic  system  there  are, 
in  each  case,  three  elements  between  the  two  divisions  of  the 
shorter  periods.  Thus  between  the  short  periods ,  potassium- 
manganese  and  copper-bromine,  we  have  the  very  important 
elements  iron,  cobalt  and  nickel.  Of  the  six  other  metals 
of  Group  VIII,  platinum  is  the  most  interesting  and  im- 
portant. All  of  the  metals  of  the  group  have  high  melting 
points — above  1500°.  The  metals  in  each  set  of  three  are 
closely  related  to  each  other  in  physical  and  chemical  prop- 
erties. Thus  iron,  cobalt  and  nickel  are  magnetic  and  easily 
soluble  in  the  ordinary  acids,  especially  nitric  acid.  The 
platinum  metals  do  not  dissolve  in  nitric  or  hydrochloric 
acid  but  are  soluble  in  aqua  regia.  Valences  of  two,  three 
and  four  are  the  most  common  in  the  group  but  higher 
valences  are  also  found. 

Iron.  Importance,  History. — Iron  was  not  discovered  till 
comparatively  late  in  the  history  of  our  race,  though  still 
before  the  beginning  of  written  history.  During  long  ages 
other  metals  were  more  used  than  iron,  but  for  several 
centuries  past  iron  has  been  more  important  than  all  other 
metals  put  together.  This  importance  is  due  chiefly  to 
three  causes: 

1.  Iron  ores  are  found  in  large  quantities  in  many  different 
parts  of  the  world  and  are  easily  mined  and  cheap. 

2.  Iron  ores  may  be  reduced  to  metallic  iron  easily  and 
cheaply  and  on  a  gigantic  scale  in  the  modern  blast  furnaces. 

323 


324  IRON,  COBALT    NICKEL,  PLATINUM 

3.  Small  amounts  of  carbon,  silicon  and  other  elements 
in  iron,  and  different  methods  of  treating  the  metal,  render 
it  hard  or  soft,  brittle  or  malleable,  and  give  products  which 
are  suitable  for  a  great  variety  of  uses. 

Occurrence  and  Metallurgy  of  Iron. — The  ores  of  iron 
which  are  used  in  its  manufacture  are  all  oxides,  or  com- 
pounds which  are  converted  into  oxides  by  heat.  The  most 
important  are  hematite,  Fe2O3,  limonite,  Fe2O3.Fe2(OH)6, 
magnetite,  Fe304,  and  siderite,  FeCO3.  The  last  is  a  con- 
stituent of  "clay  iron  stone"  which  has  been  a  very  im- 
portant ore  in  England,  the  country  which  manufactured 
more  iron  than  all  the  rest  of  the  world  during  a  large 
part  of  the  nineteenth  century. 

Metallic  iron  is  found  in  many  meteors  and  there  is  some 
probability  that  there  is  a  large  amount  of  iron  in  the  center 
of  the  earth. 

Before  the  days  of  written  history  men  discovered  how 
to  reduce  iron  ores  and  make  a  sort  of  wrought  iron  or 
steel  in  furnaces  constructed  somewhat  after  the  principle 
of  a  blacksmith's  forge.  Simple  methods  of  this  sort  have 
continued  in  use  among  primitive  people  in  India  and  Africa 
up  to  modern  times.  The  amount  of  iron  manufactured 
in  this  way  was  always  small  and  during  the  Middle  Ages 
iron  and  steel  were  comparatively  scarce  and  expensive. 

Blast  Furnaces. — About  1500  a  new  method  of  manu- 
facture similar  to  the  modern  blast  furnace  was  invented, 
but  no  one  knows  the  name  of  the  inventor.  The  old 
processes  gave  a  malleable  iron  or  steel  which  was  forged 
while  hot  but  which  was  not  melted. 

The  blast  furnace  gives  a  hard,  brittle,  easily  fusible 
iron  containing  carbon  and  called  pig  iron.  This  is  used 
for  castings  and  as  the  starting  point  for  the  manufacture 
of  all  forms  of  iron  and  steel. 

In  using  the  blast  furnace  (Fig.  50)  a  mixture  of  coke, 
iron  ore  and  limestone  is  charged  into  the  furnace  at  short 


BLAST  FURNACE 


325 


Intervals,  through  the  top,  while  a  strong  blast  of  air  is 
blown  in  at  the  bottom.  As  the  materials  make  their  way 
downward  the  coke  is 
finally  completely  burned 
at  the  bottom  of  the 
furnace,  giving  at  the  high 
temperature  and  in  the  re- 
ducing atmosphere  of  this 
part  of  the  furnace  carbon 
monoxide,  CO,  and  nitro- 
gen. Higher  up  in  the 
furnace  the  carbon  monox- 
ide reduces  the  iron  oxide 
to  metallic  iron : 

Fe203-f3CO  <=*  2Fe+3CO2 

At  the  bottom  of  the 
furnace,  where  the  com- 
bustion of  the  coke  is  com- 
pleted and  the  temperature 
is  highest,  the  iron  com- 
bines with  carbon  and 
silicon  and  melts  to  a 
liquid,  which  is  drawn  off 
from  the  hearth  from  time 
to  time.  The  impurities 
of  the  ore,  which  usually 
consist  of  silica,  SiO2,  and 
a  mixture  of  silicates,  com- 
bine with  the  calcium  oxide 
from  the  limestone  to  form 
a  liquid  slag  that  floats  on 
top  of  the  melted  iron  and  can  be  drawn  off  separately.  The 
process  goes  on  continuously,  often  for  years  at  a  time.  The 
product  of  the  furnace  is  known  as  pig  iron  or  cast  iron  It 


FIG.  50. 


326  IRON,  COBALT,  NICKEL,  PLATINUM 

may  be  used  directly  for  making  stoves  and  for  many  pur- 
poses where  the  iron  may  be  given  the  desired  form  by 
casting  it  in  molds  made  of  sand.  Cast  iron  is  brittle  and 
cannot  be  forged  or  welded. 

Wrought  Iron. — In  1784  Henry  Cort,  in  England,  invented 
a  method  of  heating  iron  and  stirring  it  in  a  current  of 
air  in  a  reverberatory  furnace.  This  was  called  the  pud- 
dling process.  The  cast  iron  is  melted  and  stirred  till  the 
carbon,  silicon,  sulfur  and  phosphorus  have  been  burned 
out,  leaving  a  pasty  mass  of  nearly  pure  iron,  which  can  be 
rolled  into  sheets,  bars  or  rods  and  afterward  forged  or 
welded  into  almost  any  form  that  is  desired.  This  process 
was  used  for  70  or  80  years  in  manufacturing  the  iron  used 
in  making  nails  and  for  a  great  variety  of  other  purposes. 
Since  the  invention  of  the  process  for  making  Bessemer 
steel,  in  1856,  the  puddling  process  has  been  gradually 
displaced  by  other  methods  which  are  more  rapid  and  less 
laborious. 

Steel. — The  most  important  property  of  steel  is  its  ability 
to  take  a  "  temper."  If  heated  to  a  moderately  high 
temperature  and  suddenly  cooled  by  quenching  it  in  water, 
it  becomes  extremely  hard.  Under  these  conditions  the 
carbon  is  uniformly  combined  with  or  dissolved  in  the 
whole  mass  of  the  iron.  If  the  steel  is  heated  and  cooled 
slowly,  the  carbon  seems  to  combine  with  a  small  part 
(about  Y^)  of  the  iron  to  form  a  definite  compound,  iron 
carbide,  Fe3C,  while  the  remainder  of  the  iron  is  free  from 
carbon  and  is  soft  and  malleable.  Intermediate  forms  of 
steel  with  varying  degrees  of  hardness  and  brittleness  may 
be  obtained  by  heating  hardened  steel  at  carefully  regulated 
temperatures. 

Bessemer  Steel. — In  1856  Henry  Bessemer,  in  England, 
described  a  process  for  manufacturing  steel  by  blowing  a 
very  strong  current  of  air  through  an  apparatus  (Fig.  51) 
containing  several  tons  of  cast  iron.  The  oxygen  of  the  air 


STEEL 


327 


burns  the  carbon  and  silicon,  of  the  iron,  and  their  combus-. 
tion  gives  so  much  heat  tkat  the  iron  remains  liquid  after 
these  elements  have  been  removed.  Wtyen  the  combustion 
of  the  carbon  and  silicon  is  complete  the  flame  at  the  mouth 
of  the  Bessemer  converter  drops  and  some*  spiegeleisen  is; 
added  to  furnish  the  amount  of  qarbpn_  required  in  the4 


FIG.  51, 

finished  steel.  After  mixing  the  contents  of  the  converter 
by  blowing  air  through  it  again  for  a  very  short  time  the 
steel  is  poured  into  ingot  molds.  When  it  has  solidified 
the  ingots  are  taken  to  the  rolls  while  still  hot  and  made  at 
once  into  rails  for  the  railways  or  into  other  forms  of  mer- 
chantable steel. 

Open-hearth  Steel. — The  original  Bessemer  process  is 
not  adapted  to  pig  iron  containing  more  than  a  very  small 


328  IRON,  COBALT,  NICKEL,  PLATINUM 

ii 


COMPOUNDS  OF  IRON  329 

amount  of  phosphorus  and  for  that  reason  and  others  a 
very  different  process  called  the  open-hearth  process  has 
grown  rapidly  in  favor  during  recent  years.  In  this  process- 
a  mixture  of  pig  iron,  ore,  steel  scraps  and  lime  or  other 
fluxing  materials  is  heated  with  gas  or  oil  in  a  furnace  con- 
nected with  regenerative  chambers  filled  with  a  checker- 
work  of  bricks  (Fig.  50).  These  chambers  are  designed 
to  absorb  heat  from  the  gases  coming  from  the  furnace 
and  afterward  to  give  up  the  heat  absorbed,  to  air  and  gas 
which  are  entering  the  furnace.  In  this  manner  a  very  high 
temperature  is  secured  with  a  comparatively  small  amount 
of  fuel. 

By  the  Bessemer  and  open-hearth  processes  iron  and  steel 
of  almost  any  desired  degree  of  hardness  and  strength  can 
be  manufactured  and  the  older  processes  of  making  wrought 
iron  and  steel  have  almost  completely  disappeared.  The 
nails  and  sheet  iron  made  by  these  processes  rust  much  more 
rapidly  than  those  made  from  the  old  puddled  iron.  It  has 
been  discovered,  however,  that  the  addition  of  a  small 
amount  of  copper  (0.2  per  cent  or  less)  greatly  lessens  this 
tendency  to  corrosion. 

Compounds  of  Iron. — Iron  forms  two  series  of  salts, 
ferrous  salts  and  ferric  salts.  In  the  ferrous  salts  such  as 
ferrous  sulfate,  FeS04,  and  ferrous  chloride,  FeCl2,  the 
iron  appears  bivalent,  while  in  ferric  salts,  as  ferric  sulfate, 
Fe2(SO4)3,  and  ferric  chloride,  FeCla,  the  iron  appears 
trivalent. 

Ferrous  Sulfate  or  Copperas,  FeSO4.7H2O,  is  easily  pre- 
pared by  dissolving  iron  or  ferrous  sulfide  in  dilute  sulfuric 
acid.  It  is  sometimes  called  green  vitriol. 

Ferrous  Hydroxide  is  a  white  precipitate  formed  by 
adding  sodium  hydroxide  to  a  solution  of  ferrous  sulfate 
entirely  free  from  oxygen  or  a  ferric  salt.  It  is  rapidly 
oxidized  by  exposure  to  the  air,  turning  green  and  then  dark 
and  finally  becoming  reddish-brown  ferric  hydroxide. 


330  IRON,  COBALT,  NICKEL,  PLATINUM 

Ferric  Sulfate,  Fe2(S04)3. — Ferrous  sulfate  may  be 
easily  oxidized  to  ferric  sulfate  by  a  great  variety  of  oxi- 
dizing agents  in  the  presence  of  sulfuric  acid.  It  gives  a 
reddish  brown  precipitate  of  ferric  hydroxide,  Fe(OH)3, 
when  sodium  hydroxide  is  added  to  its  solution. 

Ferric  Chloride,  FeCl3. — Anhydrous  ferric  chloride  is 
prepared  by  heating  iron  filings  or  turnings  in  a  current 
of  chlorine.  It  forms  green  scales  which  dissolve  readily 
in  water  to  a  yellow  or  reddish-yellow  solution.  The  solu- 
tion has  an  acid  reaction  because  the  salt  is  hydrolyzed 
by  water: 

FeCl3  +  3HOH  <=>  Fe(OH)3  +  3HC1 

If  a  solution  is  evaporated,  a  large  part  of  the  hydrochloric 
acid  escapes  and  anhydrous  ferric  chloride  cannot  be  pre- 
pared by  this  method. 

Magnetic  Oxide  of  Iron,  Fe3O4. — When  iron  is  burned 
in  oxygen  or  when  steam  is  passed  over  heated  iron  the 
magnetic  oxide  is  formed.  It  is  formed  as  a  closely  adherent 
black  coating  on  the  surface  of  red-hot  iron  exposed  to  the 
air  and  the  black  color  of  the  oxide  is  more  familiar  to  many 
people  than  the  true  color  of  iron.  The  magnetic  oxide  may 
rust  to  ferric  oxide  or  hydroxide  under  the  action  of  air  and 
water  but  it  does  so  less  easily  than  a  surface  of  metallic 
iron.  The  mineral  magnetite  has  the  same  composition 
as  the  magnetic  oxide  and  is  one  of  the  valuable  ores  of 
iron.  The  magnetic  oxide  and  magnetite  are  attracted  by  a 
magnet,  as  the  names  imply.  Other  common  compounds 
of  iron  are  only  slightly  magnetic.  The  mineral  is  some- 
times found  in  pieces  which  are  permanently  magnetic  and 
it  is  then  called  lodestone. 

Nickel  is  a  white  metal  which  resembles  iron  in  some  of  its 
properties.  It  takes  a  bright  polish  and  does  not  tarnish 
easily.  For  this  reason  it  is  deposited  electrolytically  as 
"nickel  plate"  on  iron  and  steel,  for  use  on  stoves,  bicycle 


SYMPATHETIC  INK  331 

handles  and  for  a  great  variety  of  practical  and  ornamental 
purposes. 

Nickel  silver  is  an  alloy  of  nickel, copper  and  zinc  used  as  a 
basis  for  silver-plated  ware. 

Nichrome  is  an  alloy  of  chromium  and  nickel  which  has  a 
high  melting  point  and  which  is  proving  very  useful  in 
laboratories  for  wire  used  in  electrical  resistance  furnaces, 
for  thermocouples,  for  triangles  to  support  crucibles  and 
for  many  other  purposes. 

Nickel  five-cent  pieces  are  made  of  an  alloy  containing 
75  per  cent  of  copper  and  25  per  cent  of  nickel. 

Nickel  is  bivalent  in  its  more  common  compounds,  such 
as  nickel  sulfate,  NiSO4,  and  nickel  chloride,  NiCl2. 

Cobalt  is  another  metal  resembling  iron.  It  is  very  little 
used  in  the  metallic  state  but  cobalt  oxide  gives  an  intense 
blue  color  to  glass  and  is  much  used  for  that  purpose. 
Cobalt  gives  pink  solutions  which  are  complementary  to 
the  green  solutions  of  nickel  salts.  Mixtures  of  the  two 
solutions  may  be  nearly  colorless. 

Anhydrous  cobalt  chloride  is  green,  however.  As  anhy- 
drous cobalt  chloride  absorbs  water  from  moist  air  but 
loses  it  easily  on  warming,  cobalt  chloride  is  used  as  a 
sympathetic  ink. 

If  the  trunk  and  limbs  of  a  tree  are  drawn  on  paper  with  a 
lead  pencil  and  the  leaves  are  sketched  with  a  solution  of 
cobalt  chloride,  the  leaves  will  be  nearly  or  quite  invisible 
in  moist  air  but  will  appear  on  warming  the  paper. 

Platinum. — A  large  part  of  the  platinum  of  the  world 
has  come  from  the  Ural  Mountains  in  Russia.  Some  of  it 
comes  from  Brazil  and  Colombia,  in  South  America,  and 
small  quantities  from  California,  British  Columbia  and 
Alaska.  The  metal  is  very  valuable  for  laboratory  uses 
because  of  its  high  melting  point  and  because  it  does  not 
tarnish  in  air  at  any  temperature  and  does  not  dissolve 
in  any  single  acid  or  base  in  common  use.  Its  coefficient 


332  IRON,  COBALT,  NICKEL,  PLATINUM 

of  expansion  is  so  nearly  the  same  as  that  of  glass  that 
platinum  wires  may  be  sealed  in  the  walls  of  eudiometers 
and  other  forms  of  laboratory  apparatus,  forming  a  gas- 
tight  joint  which  does  not  crack  or  leak.  For  this  reason 
platinum  has  been  much  used  for  the  leading-in  wires  of 
electric  light  bulbs. 

Platinum  has  been  used  in  Russia  for  coins,  but  the 
fluctuation  in  value  in  modern  times  has  made  such  use 
undesirable. 

The  very  great  increase  in  the  price  of  platinum  in  recent 
years  has  been  due  to  its  use  in  jewelry.  As  platinum  is 
very  important  as  a  catalyst  in  the  manufacture  of  sulfur 
trioxide,  and  for  the  oxidation  of  ammonia  to  nitric  acid, 
the  use  of  the  metal  for  articles  of  jewelry  should  be 
discontinued. 

Chloroplatinic  Acid.  Potassium  Chloroplatinate. — Plati- 
num dissolves  in  aqua  regia,  giving  a  solution  of  chloro- 
platinic acid,  H2PtCl6,  which  has  often  been  called  "platinic 
chloride."  The  latter  name  is  given  correctly  only  to  the 
compound  PtCl4,  which  may  be  prepared  by  heating  chloro- 
platinic  acid  in  a  current  of  chlorine. 

If  potassium  chloride  KC1,  is  added  to  a  solution  of 
chloroplatinic  acid,  a  precipitate  of  potassium  chloroplati- 
nate,  K2PtCl6,  separates.  This  compound  is  used  for  the 
detection  and  quantitative  determination  of  potassium. 
Ammonium  salts  give  a  similar  precipitate  of  ammonium 
chloroplatinate,  (NH4)2PtCl6. 

SUMMARY 

The  nine  elements  of  Group  VIII  are  situated  in  groups 
of  three  in  the  midst  of  the  long  periods,  between  the  first, 
and  second  divisions  of  each  period.  The  elements  vary 
in  valence,  with  valences  of  two,  three  and  four  most 
common. 


SUMMARY.     IRON,  ETC.  333 

Iron  is  the  most  important  metal.  The  manufacture  of 
articles  of  iron  and  steel  was  carried  on  before  the  beginning 
of  written  history.  Its  manufacture  on  a  very  large  scale 
is  comparatively  modern. 

Pig  iron  or  cast  iron  is  made  in  blast  furnaces.  It  con- 
tains carbon  and  silicon  as  necessary  ingredients,  and  sulfur 
and  phosphorus  as  impurities. 

Wrought  iron  and  steel  were  made  in  ancient  times 
directly  from  the  ores.  From  near  the  close  of  the 
eighteenth  till  after  the  middle  of  the  nineteenth  century 
wrought  iron  was  made  by  the  puddling  process. 

The  process  for  Bessemer  steel  was  invented  about  the 
'  middle  of  the  nineteenth  century,  that  for  open-hearth  steeT 
some  years  later.  In  all  of  the  processes  the  carbon,  silicon, 
phosphorus  and  sulfur  are  more  or  less  completely  burned 
out  of  the  iron  and  iron  or  steel  containing  a  very  small 
amount  of  carbon  is  produced. 

The  hardness  of  steel  is  dependent  on  the  amount  and  the 
form  of  the  carbon  which  it  contains. 

Copperas  is  prepared  by  dissolving  iron  in  dilute  sulfuric 
acid. 

Ferrous  hydroxide  is  prepared  by  precipitation.  It  is 
white  but  turns  dark  in  the  air. 

Ferric  sulfate  is  prepared  by  oxidizing  copperas  in  the 
presence  of  sulfuric  acid. 

Anhydrous  ferric  chloride  is  prepared  by  heating  iron  in 
chlorine.  It  is  hydrolyzed  by  water. 

Magnetic  oxide  of  iron  is  formed  by  heating  iron  to  a  high 
temperature  in  the  air. 

Nickel  resembles  iron  but  oxidizes  less  easily. 

Nickel  silver  contains  copper,  nickel  and  zinc. 

Nichrome  is  an  alloy  of  nickel  and  chromium.  It  has  a 
very  high  melting  point. 

Nickel  is  bivalent. 

Cobalt  also  resembles  iron. 


•334  IRON,  COBALT,  NICKEL,  PLATINUM 

Sympathetic  ink  faiay  be  made  from  a  solution  of  cobalt 
chloride. 

Platinum  is  found  native.  It  is  used  for  crucibles  and  as 
•a  catalyst  in  the  manufacture  of  sulfuric  acid  and  of  nitric 
•acid. 

The  most  common  laboratory  compounds  of  platinum  are 
'chloroplatihic  acid  and  potassium  chloroplatinate. 

EXERCISES 

1.  One  hundred  parts  of  water  dissolve  48  parts  of  copperas  at 
20°.     What  strength  of  dilute  sulfuric  acid  must  be  used  to  dis- 
;solve  ferrous  sulfide  and  give  a  solution  of  copperas  just  saturated 
•at  that  temperature? 

2.  If  sulfuric  acid  of  20  per  cent  is  saturated  with  ferrous  sulfate 
by  dissolving  iron  in  it,  what  per  cent  of  the  copperas  formed 
Should  separate  on  cooling  the  solution  to  20°? 

3.  If  platinum  is  worth  $1.75  per  gram,  how  much  is  a  gram  of 
chloroplatinic  acid  worth,  not  taking  account  of  the  labor  and 
materials  other  than  the  platinum  used  in  its  preparation? 

4.  How  much  pig  iron  containing  2.5  per  cent  of  carbon,  2.0 
per  cent  of  silicon  and  0.5  per  cent  of  other  elements  besides  iron 
can  be  made  from  one  ton  of  an  iron  ore  containing  60  per  cent 
of  iron? 

5.  Write  the  equation  for  the  oxidation  of  ferrous  sulfate  to 
ferric   sulfate  by  potassium   permanganate  in   the   presence   of 
sulfuric  acid. 

6.  How  many  milligrams  of  iron  in  the  form  of  ferrous  sulfate 
will  be  oxidized  by  1  cc.  of  the  solution  referred  to  in  4,  page  322? 


CHAPTER  XXXI 
ANALYSIS 

Qualitative  analysis  has  for  its  object  the  determination 
of  whether  certain  elements,  or  certain  groups  of  elements, 
such  as  the  ammonium  group,  NH4,  the  sulfate  group, 
SO4,  the  nitrate  group,  NOs,  and  others,  are  present  in  a 
given  substance  or  mixture.  In  most  cases  an  element  is 
separated  by  converting  it  into  some  well-known  and  easily 
recognizable  compound.  Thus  hydrogen  may  be  recog- 
nized by  converting  it  into  water,  or  silver  by  converting  it 
into  silver  chloride. 

It  would  be  impossible  within  the  scope  of  this  text-book 
to  give  full  directions  for  the  qualitative  analysis  of  mix- 
tures containing  all  of  the  more  common  elements  and 
radicals.  The  following  experiments  will,  however,  illus- 
trate the  methods  which  are  used  in  analysis. 

Separation  of  Lead,  Silver  and  Mercurous  Mercury, 
Metals  Whose  Chlorides  are  Nearly  Insoluble. — Try  the 
following  experiments  with  3  cc.  of  a  solution  of  lead 
nitrate.  Add  some  dilute  hydrochloric  acid,  allow  the  pre- 
cipitate to  settle,  pour  off  the  solution,  add  5  cc.  of  water, 
allow  to  settle,  pour  off  and  repeat  a  second  time.  Add  to 
the  precipitate  5  cc.  of  water,  boil  and  pour  off  the  solution 
into  a  clean  test-tube  and  notice  that  the  lead  chloride  which 
dissolves  crystallizes  from  the  solution  on  cooling.  Repeat 
the  treatment  with  hot  water  till  all  is  dissolved.  Add  a 
drop  of  dilute  sulfuric  acid  to  the  solution  of  lead  chloride. 

Repeat  the  same  experiments  with  a  solution  of  silver  ni- 
trate. When  convinced  that  silver  chloride  is*  practically 

335 


336  ANALYSIS 

insoluble  in  hot  water,  add  some  ammonia  to  the  precipi- 
tate. Test  the  solution  by  adding  nitric  acid. 

Repeat  the  experiments  with  a  solution  of  mercurous 
nitrate. 

Record  the  results  obtained  and  by  studying  them  devise 
a  method  for  separating  and  detecting  the  three  metals, 
when  all  are  present.  Prepare  a  solution  containing  the 
metals  and  demonstrate  their  separation. 

Separation  of  Mercuric  Mercury,  Lead,  Copper  and 
Bismuth,  Metals  Whose  Sulfides  are  Insoluble  in  Dilute 
Acids. — Take  3  cc.  of  a  solution  of  mercuric  chloride  and 
pass  hydrogen  sulfide  through  the  solution.  Filter  on  a 
small  filter  supported  in  a  funnel.  Wash  the  precipitate 
five  or  six  times  by  filling  the  filter  with  water  and  allowing 
the  water  to  run  through.  Rinse  the  precipitate  back  into 
the  test-tube,  allow  it  to  settle  and  pour  off  the  water. 
Boil  the  precipitate  with  dilute  nitric  acid,  then  add  some 
hydrochloric  acid  and  boil  again.  To  the  solution  contain- 
ing mercuric  chloride  add  a  solution  of  stannous  chloride, 
SnCl2.  The  precipitate  is  either  mercurous  chloride  or 
metallic  mercury,  according  to  the  amount  of  stannous 
chloride  used. 

Repeat  the  first  part  of  the  above  experiments,  using  a 
solution  of  lead  nitrate.  The  lead  sulfide  will  dissolve  in 
the  nitric  acid.  Add  to  the  solution  some  dilute  sulfuric 
acid,  pour  the  solution  into  a  porcelain  dish  and  evaporate 
nearly  to  dryness  but  not  after  fumes  of  sulfuric  acid  appear. 
When  cold  rinse  back  into  a  test-tube  and  notice  the  pre- 
cipitate of  lead  sulfate. 

Repeat  the  experiments  again  with  solutions  of  copper 
sulfate  and  of  bismuth  chloride.  After  the  evaporation 
with  sulfuric  acid  add  ammonia  to  the  solution.  Filter 
off  the  precipitate  of  bismuth  hydroxide,  BiOOH,  dissolve 
it  on  the  filter  by  dropping  dilute  hydrochloric  acid  over 
it  and  add  a  considerable  amount  of  water  to  the  solution 


QUANTITATIVE  ANALYSIS  337 

of  bismuth  trichloride,  Bids.  The  precipitate  is  bismuth 
oxychloride,  BiOCl. 

Devise  a  method  of  separating  the  four  metals  and  apply 
it  to  a  solution  containing  all  four. 

Detection  of  Sulfates,  Chlorides  and  Nitrates.— Test 
solutions  of  a  number  of  the  common  salts  of  the  laboratory 
with  barium  chloride  followed  by  hydrochloric  acid,  with 
silver  nitrate  followed  by  nitric  acid  and  with  copper 
turnings  and  sulfuric  acid.  More  sensitive  tests  for 
nitric  acifa  are  described  in  books  on  qualitative  analysis 
and  also  further  details  about  the  tests  for  sulfates  and 
chlorides. 

Quantitative  analysis  is  designed  to  furnish  information  as 
to  the  exact  quantity  of  substances  which  are  present  in  a 
compound  or  mixture.  For  this  purpose  the  element  or 
group  is  usually  converted  into  some  insoluble  compound 
which  may  be  collected  on  a  filter  and  washed  free  from 
all  other  substances.  The  filter  is  then  burned  in  a  porcelain 
or  platinum  crucible  and  the  compound  is  weighed.  Special 
precautions  are  often  required  to  avoid  reduction  of  the 
compound  by  the  carbon  of  the  filter  or  for  other  reasons. 

The  determination  of  silver  or  of  chlorine  as  silver  chloride 
and  the  determination  of  barium  or  of  the  sulfate  ion,  SO4, 
as  barium  sulfate  may  be  given  as  illustrations  of  such  a 
process. 


INDEX 


Absolute  temperature,  explana- 
tion, 34,  40. 

Absolute   zero,    explanation,    34. 

Acetaldehyde,  as  absorbent  of 
acetylene,  201. 

Acetic  acid,  in  vinegar,  20; 
preparation,  226,  235. 

Acetone,  as  absorbent  of  acety- 
lene, 201. 

Acetylene,  preparation ;  proper- 
ties; use  as  an  illuminating 
gas,  200,  208;  blowpipe,  use 
of,  26 ;  burners ;  generators,  201 ; 
decomposition  by  heat,  201; 
why  flame  gives  intense  light; 
why  explosive;  when  non- 
explosive,  201. 

Acid  calcium  phosphate,  prepara- 
tion, 171,  265. 

Acid  calcium  sulfite,  preparation, 
use,  117,  118. 

Acid  potassium  tartrate,  pre- 
paration, properties;  structure, 
use,  228;  action  in  baking 
powders,  229 

Acid  salts,  how  formed,  77,  79. 

Acid  sodium  carbonate,  action 
in  baking  powders,  229 

Acid  sodium  sulfate,  prepara- 
tion, 80;  why  so  named,  80; 
by-product  in  preparation  of 
nitric  acid,  144. 

Acid  sodium  sulfite,  preparation; 
properties,  use  as  germicide; 
for  preparation  of  sulfur  diox- 
ide, 117,  118;  use  in  tanning, 
318. 


Acid     sulfates,    properties,    125. 

Acid  taste,  indicates  hydrogen, 
20. 

Acids,  litmus  test;  sour  taste, 
11;  properties,  73;  mono- 
basic, dibasic,  tribasic,  defi- 
nition, 76,  118;  bibasic,  form 
acid  salts,  76;  modern  defi- 
nition, 78;  strength  of  in 
relation  to  solubility  of  sul- 
fides,  115. 

Actinium,  series  of  derivatives, 
273. 

Agate,  form  of  silicon  dioxide, 
238. 

Air,  relation  to  combustion,  6; 
determination  of  weight  of 
liter  of,  39;  weight  of  gram- 
molecular  volume,  137;  com- 
position of,  155;  proof  of  mix- 
ture in,  weights  of  the  four 
gases,  159,  161. 

Alabaster,  properties,  use,  110, 
261. 

Albumin,    occurrence,    230,    235. 

Alcohol,  preparation;  distillation, 
225;  source  of  ethylene,  199; 
use;  denatured,  226. 

Aldehyde,  by-product  in  prepa- 
ration of  chrome  alum,  317. 

Alfalfa,  fixation  of  nitrogen  by, 
138. 

Alizarin,    from     coal    tar,     233. 

Alkali-earth  metals,  why  so 
named ;  properties ;  sulfates, 
properties,  260;  list  of;  prop- 
erties of  compounds  of,  274. 


339 


340 


INDEX 


Alkali  metals,  properties  of  salts 
of,  217;  solubility  of  salts 
of,  280;  list  of;  properties; 
salts  of,  properties,  290,  305. 

Alkaloids,  occurrence;  properties; 
list  of,  232,  236. 

Allotropic  forms,  of  elements, 
16,  17;  oxygen,  15,  111;  sulfur, 
111;  of  phosphorus,  171;  of 
carbon,  187. 

Alloys,  of  bismuth,  181;  of 
tin,  description,  245;  of  tin, 
list  of,  251;  of  lead,  properties, 
248;  of  copper,  309;  of  gold 
and  silver,  separation  of  gold 
from,  313;  gold  coins  as,  314; 
of  manganese,  321;  compo- 
sition of  nickel  silver;  of 
nichrome,  331. 

Alum,  see  potassium  aluminium 
sulfate,  230;  use  in  manufac- 
ture of  matches,  172;  proper- 
ties; preparation,  use,  257;  as 
a  mordant  in  dyeing,  257;  use; 
kinds  of,  259;  chrome,  prepa- 
ration, 317. 

"Alum,"  for  purification  of 
water;  see  aluminium  sulfate, 
256. 

Aluminium,  occurrence,  2.54 ; 
preparation;  properties,  use, 
255;  use  in  welding,  256;  occurr- 
ence ;  preparation ;  properties ; 
use ;  welding  iron,  258 ;  manufac- 
ture by  electrolysis;  properties; 
use  in  Goldschmidt's  Thermite 
Process;  in  preparation  of 
metals,  278 ;  use  in  preparation 
of  chromium,  316. 

Aluminium  hydroxide^  use  in 
purification  of  water;  in  dyeing, 
257. 


Aluminium  oxide,  use  in  prepa- 
ration of  ethylene,  199;  source 
of  aluminium  by  electrolysis, 
255. 

Aluminium  sulfate,  preparation; 
"alum;"  use,  256;  for  purifica- 
tion of  water,  amount  used,  70; 
use  in  manufacture  of  alum, 
257;  use,  259. 

Amalgams,  composition;  proper- 
ties, 286. 

Amethyst,  form  of  silicon  dioxide, 
238. 

Ammonia,  in  soil,  138;  from 
organic  matter;  formed  by  aid 
of  catalyzer,  139;  properties, 
solubility,  why  alkaline,  140; 
preparation  as  gas,  141;  com- 
mercial preparation  of;  syn- 
thesis of,  142;  principles  of 
synthesis  of,  142;  from  action 
of  nitric  acid  on  iron,  145; 
from  organic  matter;  from 
manufacture  of  gas  or  coke; 
from  direct  union  of  hydrogen 
and  nitrogen,  152;  by-product 
of  coking,  189;  from  calcium 
cyanide  as  fertilizer,  266; 
use  in  Solvay  process,  detection 
in  drinking  water,  293,  301; 
use  to  dissolve  silver  chloride, 
311. 

Ammonia-Soda  Process,  293. 

Ammonium,  why  classed  as  an 
alkali  metal,  299;  solubility 
of  salts  of,  280;  occurrence,  306. 

Ammonium  bicarbonate,  use  in 
Solvay  process,  293. 

Ammonium  chloride,  prepara- 
tion, properties,  141,  301,  306. 

Ammonium  chloroplatinate 
preparation,  332. 


INDEX 


341 


Ammonium  hydroxide,  how 
formed;  valence  of  nitrogen  in, 
141;  how  formed,  152;  prepa- 
ration, action  reversible;  use, 
301;  properties,  306. 

Ammonium  nitrite,  formed  by 
lightning  flash,  138. 

Ammonium  salts,  how  prepared; 
effect  of  bases  on,  141;  prepa- 
ration, 152;  list  of,  302. 

Ammonium  sulfarsenite,  prepara- 
tion ;  use  in  qualitative  analysis, 
179. 

Ammonium  sulfantimonite,   181. 

Ammonium  sulfate,  in  fireproof- 
ing  cotton  goods,  246. 

Ammonium  sulfide,  solvent  of 
arsenic  trisulfide,  179;  solvent 
of  antimony  trisulfide,  181. 

Amorphous  sulfur,  properties, 
111. 

Amyl  acetate,  solvent  of  cellulose 
nitrates;  use  of  solution,  221. 

Analysis,  direct,  definition,  3; 
indirect,  5;  groups  in,  127; 
hydrogen  sulfide  basis  of 
groups  in,  112;  qualitative, 
335;  quantitative,  337. 

Analysis — synthesis,  explanation, 
42,  55. 

Anhydride,  definition,  151. 

Aniline  dyes,  why  so  called; 
properties,  233. 

Animal  charcoal,  preparation ; 
use,  189. 

Anions,  definition,  70;  in  com- 
plex cyanides,  217. 

Anode,  definition,  44,  55. 

Anthracite  coal,  how  formed; 
properties,  191. 

Antidotes,  vapor  of  alcohol  for 
chlorine  gas;  soda-lime  and 


sodium  thiosulfate  for  war 
masks;  charcoal  for  masks,  83; 
for  corrosive  sublimate,  287; 
for  snake  bite,  322. 

Antimonious  acid,  properties,  181. 

Antimony,  occurrence ;  prepara- 
tion; properties,  use,  179,  184; 
use  of  in  alloys,  180;  in  chlorine 
gas,  83;  oxides  of,  181,  184; 
reduction  of  sulfide  of,  277. 

Antimony  hydroxide,  properties, 
both  acid  and  base,  181,  184. 

Antimony  oxychloride,  prepara- 
tion, 181. 

Antimony  pentachloride,  prepara- 
tion, 83. 

Antimony  trichloride,  preparation, 
83 ;  preparation ;  properties, 
180;  hydrolysis  of,  184. 

Antimony  trisulfide,  occurrence, 
preparation,  properties;  solu- 
tion in  ammonium  sulfide. 
179,  181,  184. 

Antiseptic,  mercuric  chloride  as, 
287. 

Antitoxins,  how  formed;  use,  232, 
236. 

Apatite,  composition,  170. 

Appolinaris  water,  how  treated, 
68. 

A'qua  regia,  solvent  of  gold  or 
platinum,  146,  152,  314,  332. 

Aqueous  solutions,  presence  of 
hydrogen  ions  and  hydroxide 
ions,  72. 

Argon  family,  properties,  96;  in 
air,  how  discovered ;  properties, 
156,  160;  exception  in  periodic 
table,  168. 

Arsine,  preparation;  properties; 
test  of  arsenic  poisoning,  177, 
183. 


342 


INDEX 


Arsenic,  occurrence,  properties, 
177,  183. 

Arsenic  acid,  preparation,  salts 
of,  179,  183. 

Arsenic  trichloride,  preparation; 
hydrolysis  of,  178,  183. 

Arsenic  trioxide,  formation,  prop- 
erties; use,  177. 

Arsenic  trisulfide,  preparation; 
occurrence;  use;  how  dissolved, 
179,  183. 

Arsenious  acid,  properties,  salts 
of,  178. 

Artificial  silk,  from  cellulose 
nitrate,  221. 

Asbestos,    a   silicate,  239. 

-ate,  meaning  of,  85. 

Atmospheric  pressure,  at  sea 
level,  36. 

Atoms,  definition,  29;  not  parts 
by  weight,  52;  composite  na- 
ture of,  163. 

Atomic  number,  how  discovered; 
relation  to  atomic  weight,  168. 

Atomic  theory,  50. 

Atomic  volume,  definition,  169; 
how  determined,  168. 

Atomic  weights,  definition,  49; 
told  by  formula,  52;  give  pro- 
portion of  substances  in  com- 
bination, 56;  in  table  of  family 
groups,  100;  classification  of 
families  by,  107;  and  combin- 
ing volumes,  131;  use  for  com- 
position of  compounds,  132; 
how  determined,  133,  137; 
relation  to  periodic  system, 
133;  three  exceptions  in  table 
of  periodic  system;  relation 
of  atomic  number  to,  168;  de- 
termination of  for  calcium,  266; 
table  for  carbon  family,  237. 


Atropine,  use,  233. 

Avogadro's  law,  135;  comparison 

with     law     of      Dulong      and 

Petit,  268. 

Babbitt  metal,  alloy  of  lead,  248; 
composition,  180,  245. 

Bacteria,  aid  in  fixing  nitrogen; 
in  decay  of  organic  matter,  138; 
causes  of  disease,  232;  killed 
by  radium,  269;  oxidation  by, 
14. 

Baking  powders,  properties,  229; 
preparation,  235;  composition, 
cost,  259. 

Baking  soda,  see  sodium  bi- 
carbonate, 229. 

Barium,  occurrence,  properties, 
268. 

Barium  chloride,  properties; 
use  in  laboratory,  268,  275. 

Barium  peroxide,  preparation ; 
properties;  use,  268,  275. 

Barium  sulfate,  use,  110;  in- 
soluble in  water;  test  of  sul- 
furic  acid  in  solutions,  125; 
properties,  268. 

Barometers,  use  of  mercury  in, 
286. 

Bases,  give  hydroxide  ions  in 
aqueous  solutions,  74;  as  hy- 
droxyl  compounds,  78;  action 
on  ammonium  salts,  141,  152. 

Beehive  coke  ovens,  189. 

Beet  sugar,  see  cane  sugar,  222. 

Benzene,  occurrence ;  preparation ; 
use,  202,  208;  derivative  of 
coal  tar,  189. 

Bessemer  process  for  manufac- 
ture of  steel,  spiegeleisen  in, 
321;  change  in  manufacture 
by,  326 ;  when  invented ;  char- 


INDEX 


343 


acter  of,  333;  use  for  copper, 
309. 

Bessemer  steel,  manufacture, 
326;  use,  327. 

Bibasic  acids,  form  acid  salts,  76. 

Binary  compounds,  names  of,  15. 

1  'Biscuit"  of  earthenware. 

Bismuth,  occurrence;  properties; 
use;  alloys  of,  180;  compounds 
of;  basic  nitrate  of,  182,  184. 

Bismuth,  basic  nitrate,  prepara- 
tion; use,  182,  184. 

Bismuth  chloride,  properties, 
hydrolysis  of,  182. 

Bismuth  nitrate,  properties,  hy- 
drolysis, 182. 

Bismuth,  "  subnitrate,"  182. 

Bismuth  sulfide,  properties,  182. 

Bismuth  trioxide,  properties ;  how 
dissolved,  182,  184. 

Bituminous  coal,  properties,  191, 
192;  composition  of  sample  of, 
161. 

"Black  lead,"  popular  name 
for  graphite.  187. 

£lack  oxide  of  manganese,  use, 
properties,  321. 

.31ast  furnace,  production  of 
pig  iron,  method  of  operation, 
324;  when  invented;  use  of 
coal  in,  277. 

Bleaching,  by  chlorine  gas,  84, 
90;  by  sulfur  dioxide;  by 
chlorine,  116;  use  of  hydrogen 
peroxide,  296. 

Bleaching  powder,  how  prepared ; 
properties;  use,  86;  for  purifica- 
tion of  water,  70. 

Blowpipe,  construction  and  use, 
207;  oxyhydrogen,  why  tem- 
perature is  limited,  25;  acety- 
lene, use  of,  26. 


Blue    vitriol,    name    for    copper 

sulfate,  285;  preparation;  use, 

310,  314. 
Bone    black,    preparation;    use, 

189,  193. 

Borates,  how  formed,  254. 
Borax,     occurrence ;     derivation 

properties,  253;  use  in  welding, 

254;  use  as  preservative,  254; 

derivation;  use,  258;  use,  306. 
Borax  bends,   how  formed;  use 

in  laboratory,  254,  258. 
Bordeaux  mixture,  composition, 

310. 
Boric   acid,  use  as  preservative; 

as  eye-wash ;  properties ;  prepa- 

ration,  253,  254,  258. 
Boric  anhydride,  in  forming  bor- 

ates,  254. 
Boron,     properties ;     occurrence, 

253,  258. 

Boyle,  law  of,  statement ;  explana- 
tion, 36,  40. 
Brass,  alloy  of  copper  and  zinc, 

285. 

Bread,  manufacture,  229,  235. 
Brick,  257. 

Brines,  where  found,  68. 
British  thermal  units,  equivalent 

in  calories,  209. 
Bromides,  occurrence,  98. 
Bromine,     occurrence;     prepara- 
tion;   properties;    compounds, 

93,  98. 

Bronze,  composition,  245. 
Bunsen  burner,  flame  separated; 

products    of   flame,    205,    208; 

temperature     of     flame,     206. 
Burns,  by  sulfuric  acid,  remedies, 

124;  from  radium,   269. 
By-products    coke  ovens,  pnxU 

ucts  of,  189. 


344 


INDEX 


Cadmium,  properties ',  use  in  fusi- 
ble alloys,  285,  288. 

Cadmium  hydroxide,  properties, 
260. 

Calcium,  occurrence,  260. 

Calcium,  preparation;  properties; 
use  in  laboratory,  261;  deter- 
mination of  atomic  weight 
of,  266. 

Calcium  aluminate,  in  manufac- 
ture of  cement,  263. 

Calcium  bicarbonate,  presence  in 
water;  how  removed,  264; 
cause  of  temporary  hardness 
in  water,  275. 

Calcium  carbide,  preparation ; 
source  of  acetylene,  200;  prepa- 
ration, 208;  manufacture,  use, 
265;  calcium  cyanamide  from, 
266. 

Calcium  carbonate,  from  com- 
bustion, experiment,  11;  (lime- 
stone), effect  on  water,  62; 
occurrence,  260;  how  formed 
in  mortar,  262;  deposit  from 
"hard"  water,  264;  use  in 
Le  Blanc  soda  process,  292; 
properties,  294. 

Calcium  chloride,  in  apatite,  171 ; 
preparation;  properties;  use, 
264;  anhydrous,  use,  275;  anhy- 
drous, source  of  calcium, 261 ;  by- 
product of  Solvay  process,  293 

Calcium  cyanamide,  preparation ; 
use,  266. 

Calcium  fluoride,  properties,  95; 
in  apatite,  170. 

Calcium  hydroxide,  for  purifica- 
tion of  water,  70;  reaction  with 
chlorine,  86;  by-product  in 
forming  acetylene,  200;  prepa- 
ration, use,  262,  275. 


Calcium  light,  definition,  26. 

Calcium  nitrate,  in  soil,  138; 
source  of,  139. 

Calcium  oxide,  manufacture  of, 
261. 

Calcium  phosphate,  in  apatite, 
170;  occurrence,  170,  260;  in 
manufacture  of  phosphorus, 
171 ;  occurrence ;  use ;  properties, 
265;  use,  275. 

Calcium  silicate,  in  manufacture 
of  cement,  263. 

Calcium  sulfate,  use,  110;  anhy- 
drous, properties,  264;  cause 
of  permanent  hardness  in  water ; 
of  scale  in  boilers;  how  re- 
moved; properties,  264,  275. 

Calcium  sulfide,  in  Le  Blanc  soda 
process,  292. 

Calcium  sulfite,  preparation,  use, 
117,  118. 

Calculation  of  amounts  of  sub- 
stances in  a  reaction,  54;  of 
percentage  composition  from 
formula,  53. 

Calomel,  preparation;  use  solu- 
bility, 287. 

Calorie,  definition;  small  calorie, 
58;  equivalent  in  British  ther- 
mal units,  209. 

Calorimeter,  respiration, use  of,  13. 

Candle   flame,   description,    205. 

Cane  sugar  or  sucrose,  prepara- 
tion, 222. 

Cannel  coals,  properties,  use,  192. 

Canon  Diablo,  diamonds  from, 
186. 

Carbohydrates,  definition,  219, 
234. 

Carbolic  acid,  preparation;  use, 
228,  235;  use  of  benzene  in 
manufacture  of,  202. 


INDEX 


345 


Carbon,  allotropic  forms  of;  amor- 
phous, preparation,  properties, 
187,  188,  193. 

Carbon  family,  occurrence;  prop- 
erties; table  of  atomic  weights 
of,  237;  occurrence;  members 
of,  250. 

Carbon,  in  manufacture  of  phos- 
phorus, 171;  position  among 
the  elements;  properties;  oc- 
currence, 185;  gas,  preparation; 
properties,  use,  191;  how  ob- 
tained by  plants;  192;  heat 
of  combustion  of,  209;  relation 
to  living  matter,  237;  reduc- 
tion from  oxides,  238;  use  in 
Le  Blanc  soda  process,  292; 
use  in  reduction  of  oxides, 
277. 

Carbon  dioxide,  from  combustion 
experiment,  11;  product  of 
oxidation  in  body,  13;  use 
for  effervescent  waters,  68;  in 
air,  test  for,  155;  source  of,  156; 
support  of  vegetation;  amount 
in  air,  156;  in  air,  source  of; 
how  removed,  amount  of  test 
of  ventilation,  160;  in  air,  not 
directly  harmful,  amount  per- 
missible, 161;  source  of  carbon 
in  plants,  192;  in  illuminating 
gas,  203;  preparation,  212,  216; 
properties;  cause  of  efferves- 
cence in  soda  water,  213; 
laboratory  use  of  solid  form, 
214;  why  solution  is  slightly 
acid,  213;  reactions  of;  lique- 
faction of;  temperature  and 
vapor  pressure  of  solid  form, 
214;  conditions  of  escape  from 
solution;  melting  and  sub- 
liming points,  217;  in  manufac- 


ture  of  white  lead,  250;  use 
in  Solvay  process,  293. 

Carbon  disulfide,  preparation; 
properties;  use,  214,  217. 

Carbon  monoxide,  formation  of; 
preparation,  211;  from  im- 
perfect combustion  of  gasoline, 
198;  in  illuminating  gas,  203; 
in  water  gas,  203;  properties; 
why  so  poisonous,  212;  form- 
ation of;  properties,  poison- 
ous quality,  216;  from  potas- 
sium ferri cyanide,  217. 

Carbonic  acid,  effect  on  water, 
62. 

Carborundum,  preparation ;  prop-  . 
erties;  use,  238,  251,  255. 

Casein,     occurrence,     230,     235. 

Cassiterite,occurrence ;  reduction, 
241. 

Cast  iron,  manufacture ;  use ;  prop- 
erties, 325. 

Castner-Kellner  apparatus,  for 
manufacture  of  sodium  hydrox- 
ide, 294. 

Catalysis,  definition,  9;  of  de- 
composition of  potassium  chlo- 
rate, 9;  use  of  platinum  with 
sulfur  dioxide  and  oxygen,  119; 
use  of  nitrogen  dioxide  in  lead 
chamber  process,  121;  osmium 
and  uranium  in  forming  am- 
monia, 140;  why  osmium  and 
uranium  aid  in  forming  am- 
monia, 142;  use  of  platinized 
asbestos  in  preparation  of  nitric 
oxide  from  ammonia,  143; 
of  decomposition  of  oxalic  acid 
by  sulfuric,  211;  enzymes  as 
catalytic  agents,  224;  use  of 
platinum  for,  332. 

Catalyst,  definition,  9. 


346 


INDEX 


Catalytic  agent,  manganese  diox- 
ide as,  9;  copper  chloride  as, 
82;  copper  oxide  as,  307. 

Cathode,    definition,    44,    55. 

Cation,  definition,  70. 

Cavendish,  experimental  evi- 
dence of  argon,  156. 

Celluloid,  from  cellulose  nitrates; 
inflammable,  221. 

Cellulose,  occurrence,  234;  prop- 
erties; food  for  herbivorous 
animals,  219;  fibers  containing, 
220. 

Cellulose  hexanitrate :  properties ; 
preparation;  use,  220. 

Cellulose  nitrate,  preparation, 
220. 

Cement,  composition;  properties, 
275;  manufacture,  262;  com- 
position, 263;  per  cents,  of 
materials  in;  action  of,  263; 
for  glass  and  porcelain,  240. 

Chalcocite,  mineral  containing 
copper,  308. 

Chalcopyrite,  mineral  containing 
copper,  110,  308. 

Chamber  process,  see  lead  cham- 
ber process,  121. 

Charcoal,  burning  in  oxygen,  10; 
preparation ;  use ;  properties, 
kinds  of,  use  in  making  iron, 
188;  animal,  preparation;  use, 
189,  193. 

Charles,  law  of,  35,  40. 

Chemical  activity,  in  solutions,  65. 

Chemical  nomenclature,  binary 
compounds,  15;  oxygen  acids 
of  chlorine,  84. 

Chemistry,  definition,  2,  5. 

Chile  saltpeter,  see  sodium  ni- 
trate, 291 ;  source  of  nitric  acid, 
143,  152. 


Chlorates,  from  hypochlorites,  88. 

Chloric  acid,  from  potassium 
chlorate ;  decomposition,  87,  88. 

Chloride,  when  hydrolyzed,  when 
dissolved  and  ionized,  180. 

Chlorides,  how  prepared,  82. 

Chloride  of  lime,  how  prepared: 
properties,  86;  composition, 
275. 

Chlorides  of  phosphorus,  hydro- 
lysis of,  176. 

Chlorine,  antidotes  for,  83. 

Chlorine,  liquid,  for  purification 
of  water;  amount  used,  70; 
properties,  83,  90;  by  electro- 
lysis of  sodium  chloride,  73, 
295;  oxygen  acids  of,  84;  salts 
of  oxygen  acids  of,  85;  reaction 
with  calcium  hydroxide,  86; 
gas,  use  in  bleaching,  84,  90, 
116;  from  chloric  acid,  88; 
for  treatment  of  scrap  tin,  246; 
source  of,  291. 

Chlorine  dioxide,  from  chloric 
acid,  88;  use  in  matches,  89. 

Chloroauric  acid,  315. 

Chloroplatmic  acid,  preparation, 
332. 

Choke-damp,  name  for  carbon 
dioxide,  213. 

Cholera,  from  impure  water,  64. 

Chrome  alum,  preparation,  317, 
319. 

Chrome  green,  composition,  319. 

Chrome  iron  ore,  properties,  316. 

Chrome  steel,  use,  319. 

Chrome  tanning,  use  of  potas- 
sium dichromate  in,  317,  318. 

Chrome  yellow,  properties;  use, 
preparation,  250,  252,  317. 

Chromic  anhydride,  preparation; 
properties,  317. 


INDEX 


347 


Chromic  hydroxide,  preparation; 
properties,  318;  presence  in 
chrome  leather,  318. 

Chromic  oxide,  in  preparation  of 
chromium,  278,  316;  prepara- 
tion; properties,  318. 

Chromite,  see  chrome  iron  ore, 
316. 

Chromium,  occurrence,  prepara- 
tion, properties,  use,  316;  prepa- 
ration, 278;  use  in  nichrome, 
331. 

Cinnabar,  mineral  source  of  mer- 
cury, 285. 

Citric  acid,  in  lemons,  20. 

Clay,  composition  of,  formation, 
255;  when  plastic;  use,  241; 
composition;  fusion;  products 
of;  manufacture  of  brick,  earth- 
'enware,  porcelain,  257;  use  in 
manufacture  of  brick,  earthen- 
ware, porcelain,  239. 

Clay  iron  stone,  occurrence,  324. 

Coal,  how  formed;  varieties; 
properties,  191 ;  cannel,  coking 
and  non-coking,  192;  bitu- 
minous, properties;  anthracite, 
how  formed,  properties,  191; 
table  of  changes  in,  192;  bitu- 
minous, kinds,  193. 

Coal  tar,  from  coking  retorts, 
products  derived  from,  189; 
products  of,  208. 

Coal    tar    dyes,    properties,  234. 

Cobalt,  exception  in  periodic 
table,  168;  properties;  com- 
pounds of,  331. 

Cobalt  chloride,  properties;  use 
331. 

Cobalt  oxide,  properties;  use, 
331. 

Cocaine,  use,  233. 


Coke,  preparation;  use,  189,  193. 
preparation;  use,  189;  use 
in  blast  furnace,  324. 

Coke  ovens,  beehive,  189;  to 
recover  by-products,  189. 

Coking  coals,  properties,  use, 
192. 

Collection  of  gases,  9. 

Collision  between  elastic  bodies, 
law  of  reaction,  32. 

Collodion,  solution  of  nitrate  of 
cellulose,  221. 

Colloidal  solutions,  description; 
preparation,  240;  description; 
preparation,  251. 

Colloids,  description;  separation 
in  digestion;  effect  upon  clay, 
241. 

Combining  proportions,  law  of, 
48,  55,  132. 

Combining  volumes  and  atomic 
weights,  131. 

Combustion,  explanation  of,  6; 
quantity  of  heat  of,  13;  relation 
of  air  to,  6;  spontaneous,  how 
caused:  how  avoided,  14;  effi- 
ciency in  boilers,  194;  table  of 
heats  of  combustion  for  illumi- 
nating gas,  208. 

Combustion,  heat  of,  definition. 
17;  for  acetylene  compared 
with  carbon  and  hydrogen,  201 ; 
for  carbon  and  methane,  210. 

Complex  cyanides,  215. 

Compounds,  definition,  4,  5; 
in  terms  of  atomic  weights, 
132;  formation  of  volatile,  in- 
soluble and  slightly  ionized, 
279. 

Concentrated  sulfuric  acid,  cau- 
tion in  use  of,  124;  use  in 
preparation  of  ethylene,  199. 


-348 


INDEX 


Conservation  of  energy,  law  of,  14. 

Constant  proportion,  law  of, 
3,  48,  55. 

"Contact  process,"  in  manufac- 
ture of  sulfur  trioxide  and  sul- 
furic  acid,  119.  . 

Convection,  in  water,  60. 

Copper,  use  in  preparation  of 
sulfur  dioxide,  116;  specific 
gravity,  169;  alloys  of,  180; 
reduction  of  sulfide  of;  refining 
by  electrolysis,  278;  occurrence; 
preparation,  minerals  contain- 
ing copper,  308;  impurities 
in;  electrolytic  reduction;  prop- 
erties; use;  conductivity,  how 
impaired,  salts  of,  309;  per- 
centage in  U.  S.  silver  coins, 
311;  properties;  compounds  ofr 
preparation,  313;  percentage 
in  nickel  coins,  331;  use  in 
nickel  silver,  331. 

Copper  acetate,  in  Paris  green, 
179. 

Copper  arsenite,  in  Paris  green, 
179. 

Copper  chloride,  as  catalytic 
agent,  82. 

Copper  Matte,  mixture  of  sulfides 
from  sulfide  ores;  properties; 
reduction,  309. 

Copper  oxide,  reduction,  24,  26; 
use  in  determining  composition 
of  water,  45;  in  preparation  of 
sulfur  dioxide,  116;  as  catalytic 
agent,  307. 

Copper  pyrites,  see  chalcopyrite 
308. 

Copper  sulfate,  by-product  in 
preparation  of  sulfur  dioxide, 
116;  popular  name  of,  285; 
preparation;  use,  310. 


Copperas,  preparation,  329,  333. 

Corrosive  sublimate,  solubility; 
antidote;  use,  287,  288;  prepa- 
ration, 289. 

Corundum,  occurrence;  use,  255. 

Cream  of  tartar,  see  acid  potas- 
sium tartrate,  229. 

"Creosote"  for  treating  lumber, 
189. 

Cryolite,  in  electrolysis  of  alumin- 
ium oxide,  255. 

Crystallization,  purification  by,  2. 

Crystalloids,  description,  241. 

Cubic    centimeter,    defined,    28. 

Cupric  chloride,  preparation,  83. 

Cupric  oxide,  in  preparation  of 
sulfur  dioxide,  116. 

Cuprous  chloride,  preparation,  83; 
properties,  310;  preparation; 
solubility,  314. 

Cuprous  oxide,  preparation;  test 
for  glucose,  309,  314. 

Cuprous  sulfide,.  in  copper 
matte,  309. 

Cyanides,  poisonous  quality,  217; 
soluble  as  complex  cyanides, 
215. 

Cyanide  process,  in  gold  mining, 
312. 

Davy  safety  lamp,  construction 
of,  why  safe,  197;  why  in- 
vented, 208. 

Davy,  Sir  Humphrey,  preparation 
of  potassium  and  sodium  by 
electrolysis,  278. 

Deliquescence,    definition,  69. 

Denatured  alcohol,  preparation; 
use,  226. 

Developer  in  photography,  312. 

Dextrin,  preparation;  use,  224, 
234. 


INDEX 


349 


Di-,  prefix,  174. 

Diabetes,  change  of  sugar  in,  223. 

Dialysis,  definition,  241;  use. 
251. 

Diamonds,occurrence;  properties; 
how  formed,  185,  192;  arti- 
ficial, 186;  use,  187. 

Diastase,  how  formed ;  properties, 
224. 

Dibasic  acids,  definition,  76,  118. 

Diffusion  of  gases,  experiment, 
29;  explained  by  kinetic  theory, 
31. 

Direct   analysis,    definition,  3,  5. 

Disease  germs  in  water,  64. 

Disinfectants,  sulfur  dioxide; 
formaldehyde,  115;  zinc  chlo- 
ride, 288;  potassium  perman- 
ganate asr  322. 

Disodium  phosphate,  ionization 
of,  176;  product  in  digestion; 
function  in  body,  291. 

Dissociation,  defined,  150,  153; 
of  water,  25;  of  nitrogen 
tetroxide,  150. 

Distillation,     purification    by,  2. 

Divisions  of  Group  I,  contrasts 
between  elements  of,  308. 

Divisions  of  Group  VI,  contrasts 
between  elements  of,  316,  318. 

Dolomite,  composition,  261;  com- 
position, 283. 

Drummond  light,   definition,  26. 

"Dry  plates,"  preparation;  use 
311. 

Dulong  and  Petit,  law  of,  for 
exceptional  elements,  137;  law 
of,  comparison  with  law  of 
Avogadro,  267 ;  table  for  appli- 
cation of,  268. 

Dust,  explosive  mixtures  with  air, 
198. 


Dyeing,  potassium  dichromate 
as  mordant  in,  317. 

Dyes,  natural,  source,  list  of; 
artificial,  aniline;  properties, 
233;  use  of  hydrocarbons  in 
manufacture  of,  202;  vege- 
table and  artificial,  236. 

Dynamite,  from  nitroglycerine, 
125;  manufacture,  how  ex- 
ploded, 228,  235. 

Earthenware,  257. 

Efflorescence,   definition,  69. 

Eggs,  preservation,  use  of  sodium 
silicate,  240,  296. 

Elastic  bodies,  law  of  rebound, 
32,  40. 

Electric  arc,  temperature  of,  25; 
use  in  manufacture  of  nitric 
oxide  and  nitric  acid,  149. 

Electric  current,  between  metals, 
cause  of,  242;  why  direction  of 
current  is  wrongly  stated,  245. 

Electrodes,  definition,  43,  55;  gas 
carbon  for,  191. 

Electrolysis,  definition,  43;  of 
sulfuric  acid,  42,  55;  explana- 
tion, 65;  of  sodium  chloride, 
fused;  in  aqueous  solution, 
72,  78;  in  manufacture  of 
aluminium,  255 ;  in  preparation 
of  calcium,  261;  of  salt  solu- 
tion, Castner-Kellner  apparatus 
for,  294;  of  salt  solution,  305. 

Electrolyte,  definition,  44;  separa- 
tion into  ions,  65. 

Electrolytic  methods,  for  refining 
copper;  for  manufacture  of 
aluminium,  sodium  and  mag- 
nesium, 278. 

Electromotive  series,  definition, 
251;  explanation;  table  of^  243;. 


350 


INDEX 


to  determine  position  of  metals 
in,  244. 

Electrons,  definition;  constituent 
part  of  atoms  of  some  elements; 
weight  of,  163;  relation  to 
potential  between  metals;  to 
electric  current  between  metals, 
242;  in  radium  disintegration, 
272. 

Electropositive,    definition,    243. 

Electroscope,  gold  leaf,  to  detect 
radium,  269. 

Elements,  definition,  4;  example, 
number  of,  4,  5;  most  com- 
mon, 4;  groups  of,  92;  families 
of,  properties,  table,  100;  rate 
of  combination  of,  dependent 
on  temperature,  120;  disinte- 
gration of,  270;  how  detected 
in  stars,  303. 

Emery,  see  corundum,  255. 

Endothermic  compounds,  reac- 
tions, definition,  201. 

Energy,  kinetic;  potential,  2; 
law  of  conservation  of,  14; 
indestructibility,  17;  produc- 
tion of,  from  food,  230;  from 
radium,  269. 

Enzymes,  description,  224. 

Equations,  explanation ;  writing 
of,  52;  what  they  represent, 
56;  how  combined,  86. 

Equilibrium,  between  water  and 
vapor,  69;  in  reactions,  74; 
in  preparation  of  sulfur  triox- 
ide,  118;  in  reversible  reaction, 
effect  of  temperature  on, 
121. 

Etching  glass,  by  hydrofluoric 
acid,  96. 

Ethyl  alcohol,  preparation,  224, 
235. 


Ethylene,  preparation ;  occur- 
rence; compounds;  why  termed 
unsaturated ;  gives  luminous 
quality  to  gas,  199,  208. 

Ethylene  bromide,  preparation, 
199. 

Ethylene  chloride,  preparation, 
199. 

Explosive  mixtures,  composition, 
197. 

Explosives,  dependence  on  nitric 
acid,  143,  152. 

False  gold  leaf,  composition,  83. 

Fats,  substances  composed  of; 
acids  in;  salts  derived  from, 
227;  acids  in;  saponification, 
235. 

Fatty  acids,  salts  of,  227. 

Ferric  chloride,  preparation;  prop- 
erties, 330;  anhydrous,  prepa- 
ration, 333. 

Ferric  hydroxide,  preparation, 
330. 

Ferric  oxide,  use  in  welding,  256; 
use  in  Goldschmidt's  Thermite 
process,  278. 

Ferric  sulfate,  preparation,  330, 
333. 

Ferrocyanogen  ion,  description, 
216. 

Ferro-manganese,  composition 
use;  preparation,  321. 

Ferrosilicon,  manufacture;  prop- 
erties, 238. 

Ferrous  chloride,  how  formed,  21. 

Ferrous  cyanide,  in  preparation 
of  potassium  ferrocyanide,  215. 

Ferrous  hydroxide,  preparation, 
properties,  329,  333. 

Ferrous  silicate,  in  reduction  of 
copper.  309. 


INDEX 


351 


Ferrous  sulfate,  how  formed,  21; 
electrolysis  of,  216;  popular 
name  of,  green  vitriol,  285; 
preparation,  329. 

Ferrous  sulfide,  precipitation  of, 
115;  hydrogen  sulfide  from,  112; 
in  copper  matte,  309. 

Fertilizers,  from  mineral  phos- 
phates, use  of  sulfuric  acid,  125; 
use  of  gypsum,  261;  forms  of 
phosphates  used  as,  265;  use 
of  ammonia  from  calcium 
cyanide,  266;  forms  of  phos- 
phates used  as,  275;  use  of 
sodium  nitrate  as,  291;  potash 
salts  in,  298. 

Fire,  6;  early  methods  of  get- 
ting, 172. 

Firedamp,  popular  name  for 
methane;  cause  of  explosions 
in  mines,  197. 

Fireproofing  fabrics,  with  soluble 
glass,  ^40;  of  cotton  goods  with 
sodium  stannate,  246;  use  of 
sodium  silicate,  296. 

Flames,  what  luminosity  de- 
pends upon,  205;  temperature 
of,  206;  blowpipe,  oxid  zing, 
reducing,  207,  209;  what  lum- 
inosity depends  upon,  208; 
colored,  how  produced,  302, 
306. 

Flash  light  powders,  composition, 
283. 

Flashing  point,  of  gasoline,  of 
kerosene,  208. 

Flint  glass,  composition,  239, 

Flowers  of  sulfur,  definition,  110. 

Fluorine,  occurrence,  properties; 
how  prepared,  95,  98;  prepara- 
tion, 99. 
Fluorite,  use  as  flux,  261. 


Fluorspar,  properties,  95;  source 
of  fluorine;  use  as  flux,  201. 

Food,  production  of  heat  and 
energy  from,  230,  235;  for 
plants,  sources  of,  substances, 
231 ;  formation  of  tissues  from, 
231,  235. 

Formulas  for  molecules,  51 ;  what 
they  designate,  56;  graphical, 
103;  represent  gram  molecular 
volume,  135;  of  elementary 
gases,  how  determined,  136; 
for  nitrogen  tetroxide  and 
nitrogen  dioxide  at  different 
temperatures,  151  ;  for  va- 
lences in  different  groups, 
166. 

Frasch,  Hermann,  process  for 
getting  sulfur  in  Louisiana,  109. 

Fraunhofer  lines,  discovery  of, 
explanation  of,  303,  307. 

Fructose,  how  prepared;  prop- 
erties, 223. 

Furnace,  Bessemer,  327;  open 
hearth,  328;  reverberatory,  use 
in  manufacture  of  wrought 
iron,  326. 

Fusible  alloys,  use  of  cadmium  in, 
288. 

Galena,  properties,  247;  reduc- 
tion of,  310. 

Galvanized  iron,  how  made;  when 
poisonous,  284. 

Garnet,  a  silicate,  239. 

Gas,  determination  of  weight  of, 
38,  40;  measure  of  pressure  of, 
40;  law  of  partial  pressures, 
61 ;  illuminating,  properties, 
203;  water,  properties,  203; 
producer,  manufacture,  use, 
204. 


352 


INDEX 


Gases,  collection  and  storage  of, 
9;  illustration  of  diffusion,  29; 
kinetic  theory  of,  31;  diffusion 
explained  by  kinetic  theory,  31 ; 
law  of  change  of  volume  with 
temperature,  35;  law  of  change 
of  volume  with  pressure,  37; 
reduction  to  standard  volume, 
37;  diffusion  explained  by 
kinetic  theory,  40;  combina- 
tion by  volume,  diagram,  130; 
combination  by  volume,  Gay 
Lussac's  law,  131,  136;  "per- 
fect," 132;  why  "noble,"  so- 
called,  158. 

Gas  Carbon,  preparation ;  proper- 
ties, use,  191,  193. 

Gasoline,  composition;  use,  198. 

Gasometer,    method    of   use,    9. 

Gastric  juice,  dependence  on 
salt,  291. 

Gay  Lussac's  law  of  combination 
by  volume,  131. 

Gay-Lussac  tower,  use  in  manu- 
facture of  sulfuric  acid,  123. 

Germicide,  acid  sodium  sulfite 
as,  118;  use  of  sulfur  dioxide 
as,  115;  use  of  potassium  per- 
manganate, 322. 

Germs,  carried  by  water,  how 
removed,  64,  69. 

Glass,  composition;  manu- 
facture; properties;  use; 
varieties,  239,  251;  etching  by 
hydrofluoric  acid,  96. 

Glazes,  for  clay  products,  258, 
259. 

Glover  tower,  use  in  manufacture 
of  sulfuric  acid,  123. 

Glucose,  how  prepared;  use;  in 
diabetes,  223,  234;  use  in  pre- 
paring cuprous  oxide-  309. 


Gluten,    preparation,    230,    235. 

Glycerine,  see  glycerol,  227. 

Glycerol,  from  fats,  228,  235; 
manufacture;  use,  227,  235. 

Glyceryl,     in     fats,     227,     235. 

Gold,  action  of  aqua  regia  upon, 
146;  occurrence,  separation, 
312,  314;  percentage  in  coins, 
properties,  oxides  of;  chlorides 
of,  313. 

Goldschmidt's  thermite  process, 
use,  256. 

Gram,  defined,  28. 

Gram-atom,  how  designated,  56; 
simple  numbers  in  unit  volume, 
132;  weight  of,  137. 

Gram-molecule,  how  designated, 
56;  weight  of,  volume  of,  134. 

Gram  molecular  volume,  meas- 
ure of,  134,  136. 

Granite,  composition,  239. 

Graphical  formulas,  103;  to  ex- 
press valence,  107. 

Graphite,  occurrence,  prepara- 
tion; use,  properties,  187,  193. 

Gravitation,  law  of,  1. 

Green  vitriol,  see  ferrous  sulfate, 
285. 

Group  zero,  155. 

Group     I,     first    division,     290. 

Group    I,    second    division,    308. 

Group     II,     first    division,     260. 

Group  II,  second  division,   283. 

Group    III,  253. 

Group  IV,  185,  237. 

Group  V,  138,  170. 

Group  VI,  109. 

Group  VI,  second  division,  316. 

Group  VII,  92. 

Group  VII,  second  division,  320. 

Group  VIII,  characteristics  of; 
valence,  323. 


INDEX 


353 


Groups  of  elements,  92 ;  of  metals 
in  qualitative  analysis,  112, 
127;  of  periodic  system,  162. 

Gun  cotton,  manufacture  by  sul- 
fur ic  acid,  125;  preparation; 
properties,  220. 

Gunpowder,  composition,  cause 
of  explosion,  how  rate  of  burn- 
ing is  regulated,  300,  306; 
preparation  of  saltpeter  for, 279. 

Gypsum,  use,  110;  properties, 
use,  261;  in  manufacture  of 
plaster  of  Paris,  264. 

Half-life    period    of    radioactive 

elements,  276. 
Half-metals,  distinguished  from 

non-metals  and  metals,  237. 
Halogen     family,     list;     atomic 

weights,   92,  98;  properties  of 

compound  of,  97,  99. 
<:Hard  finish,"  use  of  plaster  of 

Paris    in,  264. 

"Hard"  water,  explanation,  263. 
Heat   of  combustion,   definition, 

13,  17. 
Heat,   production  of,  from  food, 

230. 
Helium,  boiling  point,  35;  where 

found;  how  formed,  158,  160; 

constituent   part  of   atoms   of 

some     elements;     weight     of, 

163;      by     decomposition    of 

radium;    cause   of   radioactive 

phenomena,  270. 
Hematite,  ore  of  iron,  324. 
Henrys  law  of  gas  pressure,  61; 

of  absorption  of  gas  by  water, 

159;  applied  to  carbon  dioxide, 

213. 
High-speed   tools,    tungsten    in, 

318. 

23 


Hunyadi  water,  mineral  content 
of,  283. 

Hydrates,  definition,  68. 

Hydraulic  cement,  see  cement, 
263. 

Hydraulic   mining   of   gold,  312. 

Hydriodic  acid,  use  in  prepara- 
tion of  phosphonium  iodide, 
173. 

Hydrocarbons,  number  of;  marsh 
gas  series  of,  195;  structure 
of  methane  series  of;  graphic 
formulas,  196;  list  of  series  and 
compounds,  202;  valences  of 
carbon  and  hydrogen  in,  207. 

Hydrochloric  acid,  with  zinc; 
with  iron,  21;  how  prepared; 
properties,  80,  90;  boiling  point 
of  solutions,  81,  90;  action  of 
oxidizing  agents  on,  82;  action 
on  metals,  81;  action  on 
hydroxides  and  oxides,  82,  90; 
standard  for  unit  volume,  132; 
by  hydrolysis  of  chlorides  of 
phosphorus,  176;  source  in 
.digestion,  291. 

Hydrocyanic  acid,  preparation ; 
use;  poisonous  quality,  215, 
217. 

Hydroferrocyanic  acid,  relation 
to  potassium  ferrocyanide; 
preparation  of  hydrocyanic 
acid  from,  215. 

Hydrofluoric  acid,  how  prepared; 
use  in  etching  glass,  96,  99. 

Hydrogen,  by  passing  steam  over 
red  hot  iron;  by  action  of 
sodium  or  potassium  on  water, 
19;  in  substances  of  acid  taste, 
20;  from  acids  and  metals,  21; 
danger  of  explosion;  proper- 
ties, 21;  properties;  explosive 


354 


INDEX 


mixtures;  occurrence;  reduc- 
tion by,  22;  liquid,  solid,  23; 
by  action  of  sodium  or  potas- 
sium on  water,  26;  properties; 
occurrence,  26;  velocity  of 
molecules,  33;  diffusion  ex- 
plained by  kinetic  theory,  33; 
as  illustration  of  absolute 
temperatures,  34;  liquefying 
temperature,  34;  weight  of  liter 
of,  43;  ions,  in  aqueous  solu- 
tions, 72;  ions,  from  acids,  73; 
replaced  by  metals,  73,  78; 
weight  of  liter  of,  99;  valence 
of,  105. 

Hydrogen  ions,  from  nitric  acid, 
114- 

Hydrogen  peroxide,  properties, 
49;  preparation  from  barium 
peroxide,  275;  preparation,  use 
in  bleaching,  296. 

Hydrogen  sulfide,  preparation, 
properties;  occurrence;  use  in 
chemical  analysis,  112,  127; 
occurrence;  use  in  chemical 
analysis,  127;  weight  of  liter, 
128;  in  illuminating  gas,  203; 
in  preparation  of  zinc  sulfide, 
285. 

Hydrolysis,  definition,  176,  183; 
of  salts  of  phosphoric  acid,  175; 
of  chlorides  of  phosphorus,  176; 
of  arsenic  trichloride,  178; 
of  bismuth  compounds,  182; 
of  salts  of  phosphoric  acid,  183; 
of  salt  in  digestion,  291. 

Hydrosulfides,  preparation,  115, 
127. 

Hydroxide  ions,  in  aqueous  solu- 
tion, 72;  from  bases,  74. 

Hydroxyl  compounds,  how  ion- 
ized, 78. 


Hypo-,  meaning  of,  85. 

Hypochlorites,  preparation,  86. 

Hypochlorous  acid,  formed  in 
bleaching,  84. 

Hypophosphorous  acid,  mono- 
basic, 176,  183. 

"Hyposulfite,"  sodium,  126. 

-ic,  meaning  of,  84. 

Ice,  latent  heat  of  fusion  of,  58,  69. 

-ide,  ending  for  binary  com- 
pounds, 15. 

Illuminating  gas,  manufacture ; 
composition;  202,  208;  what  il- 
luminating quality  depends 
upon,  203;  table  of  heats  of 
combustion  for,  209;  composi- 
tion of  different  kinds,  210. 

Indestructibility  of  matter,  law 
of,  12. 

India  rubber,  use  of  sulfur,  110. 

Indicators,  definition ;  use,  75,  79. 

Indigo,     artificial     product,  233. 

Indirect  analysis,  5. 

Insoluble  compound,  conditions 
for  formation  of,  279. 

Invert  sugar,  how  prepared,  223, 
234 ;  why  so  named ;  occurrence, 
223. 

Iodine,  occurrence;  preparation; 
properties,  94,  98;  tincture  of, 
preparation;  use,  94,  98;  com- 
pounds; properties,  98;  in  thy- 
roid gland,  95,  98;  test  for 
starch,  222. 

lonization,  in  solutions,  65;  of 
acids  in  aqueous  solutions,  73; 
of  water,  74;in  dilute  solutions, 
77;  of  nitric  acid,  114;  of 
phosphoric  acid,  175;  of  diso- 
dium  phosphate,  176;  of  air  by 
radium,  269. 


INDEX 


355 


Ions,  in  solutions,  65;  positive- 
negative;  migration  of,  66; 
in  formation  of  precipitates; 
formation  reversible,  113;ferro- 
cyanogen,  216;  transfer  from 
metals  to  solution,  242. 

Iron,  burning  in  oxygen,  11;  used 
to  decompose  water,  19;  his- 
tory of  use;  blast  furnace;  im- 
portance of,  277;  reduction  of 
oxides  of,  277;  importance; 
history,  323,  333;  occurrence; 
manufacture;  ores  of,  324; 
pig  or  cast,  manufacture,  use, 
324;  wrought,  manufacture, 
use,  326;  salts  of,  329. 

Iron  carbide,  in  manufacture  of 
steel,  326. 

Iron  pyrites,  source  of  sulfur 
dioxide,  122;  in  manufacture  of 
sulfuric  acid,  110. 

Isotopes,  compounds  of,  excep- 
tion to  law  of  constant  pro- 
portion, 3. 

-ite,  meaning  of,  85. 

Jasper,  form  of  silicon  dioxide, 

238. 
Jelly,  manufacture,  224. 

Kaolin,  a  silicate,  239;  how 
formed-,  255. 

Kerosene,  use,  flashing  point  of, 
199. 

Kilogram,  defined,  28. 

Kimberly,  diamonds  from,    185. 

Kindling  temperature,  explana- 
tion, 12,  17. 

Kinetic  energy,  2. 

Kinetic  theory  of  gases,  explana- 
tion, 31,  40. 


Kipp  generator,  use  of,  21. 
Krypton,    where   found,  158. 

Lacquers  from  cellulose  nitrate, 
221. 

Lampblack,  preparation;  use; 
properties,  188,  193. 

Latent  heat  of  fusion  of  ice.  58, 
69;  of  vaporization  of  water, 
59,  69. 

Lavoisier's  experiment,  propor- 
tion of  oxygen  to  nitrogen  in 
air,  7. 

Law  of  gravitation,  1;  of  con- 
stant proportion,  3;  definition, 
5;  of  indestructibility  of  mat- 
ter, 12;  of  conservation  of 
energy,  14;  of  velocity  of 
molecules  in  gases,  33;  of 
change  in  volume  of  gas 
with  temperature,  35;  of 
Charles,  35,  40. 

Law  of  Boyle,  statement;  ex- 
planation, 36,  40;  of  change  of 
volume  of  gas  with  pressure, 
37;  of  multiple  proportions, 
49,  56;  of  combining  propor- 
tions, 48,  55,  132;  of  constant 
proportion,  48,  55;  partial 
pressures  of  gases  (Henry's 
law),  61;  of  vapor  pressures, 
63. 

Law,  Gay  Lussac's  for  combina- 
tion of  gases,  131;  Henry's, 
of  gas  pressure,  61;  of  van't 
Hoff-Le  Chatelier,  120,  142, 
148;  of  Avogadro,  135;  of 
Dulong  and  Petit,  for  excep- 
tional elements,  137;  of  Henry, 
159;  of  Dulong  and  Petit, 
267;  table  for  application  of, 
268. 


356 


INDEX 


Leaching,  process  described,  297. 

Lead  chamber  process  for  manu- 
facturing sulfuric  acid,  121, 
128,  247. 

Lead,  use  of  alloys  of,  180; 
danger  of  use  for  water  pipes, 
247;  oc  urrence;  preparation; 
use;  oxides  of,  salts  of,  247,  252; 
differences  in  atomic  weight, 
273;  reduction  of  sulfide  of, 
278;  use,  247;  alloys  of;  oxides 
of,  248. 

Leadacetate,preparation;use,250. 

"  Lead  burning,"  122. 

Lead  chloride,  solubility;  prepa- 
ration, 250. 

Lead  c  h  r  ornate ,  preparation ; 
properties;  use,  317;  see  chrome 
yellow,  252;  use,  319. 

Lead  dioxide,  preparation;  use 
in  storage  batteries,  248. 

Lead  nitrate,  preparation,  248, 
249. 

Lead   oxide,   preparation,  247. 

"Lead"  pencils,  made  from 
graphite,  187. 

Lead  plumbate,  see  red  lead,  252. 

Lead  sulfate,  preparation,  247; 
in  storage  batteries,  249. 

Lead  sulfide,  precipitation  of, 
114. 

Le  Blanc  soda  process,  three 
operations  of ,  291,  305. 

Legumes,  fixation  of  nitrogen  by, 
138,  152. 

Lemons,     citric     acid     in,     20. 

Light,  ultra-violet  for  purifica- 
tion of  water,  70. 

Lignite,     properties,     191,     192. 

Lime,  manufacture  of,  261,  275; 
"slaking"  of,  11;  for  purifica- 
tion of  water,  70. 


Lime  light,  definition,  26. 

"  Lime -nitrogen, "  preparation ; 
use,  266.  . 

Lime -sulfur  wash,  in  spray  for 
vegetation,  110. 

Lime  water,  preparation,  11. 

Limestone  (calcium  carbonate); 
effect  on  water,  62;  composi- 
tion, 260;  use  in  blast  furnace, 
324. 

Limonite,     ore     of     iron,  324. 

Liquid     air,     how     obtained,  10. 

Liquid  hydrogen,  properties,  22. 

Liter,  defined,  28. 

Liter  of  a  gas,  determination  of 
weight  of,  38. 

Litharge ,     preparation ,  248 . 

Lithopone,    composition,    289. 

Litmus,  affected  by  acids,  11; 
as  an  indicator,  76;  action 
of  sodium  hydroxide  on,  20. 

"Liver  of  sulfur,"  use  by  Cav- 
endish to  absorb  oxygen,  157. 

Lodestone,     properties,  330. 

Lubricant,    graphite,    187. 

Lubricating  oils,  product  from 
petroleum,  199. 

Lunar  caustic,  see  silver  nitrate, 
311. 

Lye,  how  obtained ;  use  in  making 
soft  soap,  297. 

Magnesite,  composition,  283,  330; 

properties,  316;    ore    of    iron, 

324. 
Magnesium,  manufacture  by 

electrolysis,    278;    occurrence; 

properties    of    compounds   of; 

use;  compounds  of,   283,  284, 

288. 
Magnesium  ammonium  arsenate, 

179. 


INDEX 


357 


Magnesium  carbonate,  occur- 
rence, 261 ;  in  mineral  form,  283. 

Magnesium  sulfate,  occurrence; 
properties,  283;  use  in 
manufacture  of  matches,  172. 

Magnetic  oxide  of  iron,  from 
combustion  experiment;  prop- 
erties, 11;  of  iron,  by  passing 
steam  over  iron,  19;  reduc- 
tion of,  24,  27;  preparation, 
properties,  330;  preparation, 
333. 

Malachite,  mineral  containing 
copper,  308. 

Malt,  agent  in  fermentation,  224. 

Maltose,  preparation ;  properties, 
223,  234. 

Manganese,  occurrence;  prop- 
erties; use,  320,  322;  alloys 
of,  321;  oxides  of,  321. 

Manganese  chloride,  preparation, 
83. 

Manganese  dioxide,  catalytic 
agent,  9;  to  oxidize  hydro- 
chloric acid,  83;  occurrence, 
use,  320,  322. 

Maple     sugar,    description,  222. 

Marble,    how   formed,  261. 

Marsh  gas,  popular  name  for 
methane,  196. 

Marsh  gas  series  of  hydrocar- 
bons, 195. 

Matches,  use  of  phosphorus,  172. 

Matter,  law  of  indestructibility, 
12. 

Mauve,  source  of,  233. 

Meerschaum,     a     silicate,     239. 

Meker  burner,  temperature  of 
flame,  207. 

Melting  points  of  elements,  rela- 
tion to  position  in  periodic 
system,  168. 


Mercuric  chloride,  source  of 
calomel;  properties;  antidote, 
287. 

Mercuric  fulminate,  use;  prop- 
erties, 287,  288. 

Mercuric  oxide,  for  preparation 
of  oxygen,  7 ;  use  in  laboratories, 
286. 

Mercurous  chloride,  preparation, 
use,  287,  288. 

Mercury,  occurrence ;  prepa- 
ration, properties;  use,  amal- 
gams; valence,  285,  286,  288; 
use  in  Castner-Kellner  appara- 
tus, 295. 

Mesothorium,  use  in  radiolite 
watches,  273,  276. 

Metallic  properties,  variation  in 
periods  and  groups,  166. 

Metallurgy,  definition,  277. 

Metals,  replace  hydrogen  in 
acids,  73,  78;  groups  of  in 
qualitative  analysis,  112;  prop- 
erties of  salts  of  alkali,  217; 
solution  pressure  of,  explana- 
tion; cause  of  electric  cur- 
rent between  metals,  242; 
reduction  of  ores  of,  277;  of 
second  division,  group  II, 
characteristics  of,  283,  287. 

Metaphosphoric    acid,    174. 

Metastannic  acid,  preparation, 
145;  preparation;  use  as  test 
in  analysis,  246,  251. 

Meter,  defined,  28. 

Methane,  properties,  196;  prepa- 
ration, 197;  occurrence,  prepa- 
ration; explosive  mixtures 
of,  208;  heat  of  combustion  of, 
210. 

Methane  series  of  hydrocarbons, 
structure  of,  196. 


358 


INDEX 


Metric   weights   and   measures; 

equivalents,  28,  39. 
Migration  of  ions,  66. 
Milk   of  lime,  preparation,  262. 
Millimeter,  defined,  28. 
Mineral     phosphates,     use     for 

fertilizers,  125. 
Mineral     waters,     natural     and 

artificial,   67;   mineral  content 

of,  283. 

Mispickel,  source  of  arsenic,  177. 
Mixtures,     separation     of,     2. 
Moissan,    method    of    preparing 

fluorine,     96;    preparation    of 

diamonds,  186. 
Molecular    theory,    explanation, 

29,  40. 
Molecular  weight  of  compound, 

how  determined,  134. 
Molecules,    definition,    29,     39; 

law  of  .velocity  in  gases,   33; 

number   in  cubic  centimeters, 

33,   40,   273,    276;  number  in 

equal  volume,  135;  number  in 

unit  volume,  135. 
Mono-,  prefix,  174. 
Monobasic  acid,  definition,  76, 

118. 
Morphine,    source;    use;  danger 

of  use,  233. 
Mortar,  preparation;  use;  cause 

of  strength,  262. 
Multiple     proportions,    law    of, 

49,  56. 

Names  of  binary  compounds,  15; 

of    oxygen    acids    of    chlorine 

and  salts,  84,  85. 
Naphthalene,  derivative  of  coal 

tar,  189. 

Natural    gas,    occurrence,  196. 
Natural  waters,  61. 


"Negative"  in  photography,  ex- 
planation, 312. 

Negative  and  positive  valences, 
104,  106. 

Neon,  where  found,    158. 

Neutral  solution,  definition,  75, 
79 

Neutralization,  definition,  78. 

Nichrome,  composition  of  alloy; 
use,  331,  333;  high  melting 
point,  333. 

Nickel,  properties;  use;  alloy  of, 
330;  percentage  in  coins,  331. 

"Nickel  plate,"  how  deposited; 
use,  330. 

Nickel  silver,  composition  of 
alloy;  use,  331,  333. 

Nicotine,  occurrence;  properties, 
233. 

Niton,  how  formed;  properties, 
158,  161;  how  produced  from 
radium,  270;  properties,  272. 

Nitrates,  from  decay  of  organic 
matter,  144. 

Nitrates  of  cellulose,  uses  in 
solution,  221. 

Nitric  acid,  ionization,  114;  use 
in  manufacture  of  sulfuric 
acid,  J22,  128;  sources  of,  143; 
oxidation  with,  144;  prepara- 
tion, properties,  144;  action 
on  metals,  145;  preparation, 
use,  as  oxidizing  agent,  152; 
source  of;  use,  291. 

Nitric  oxide,  formed  by  electricity, 
139;  preparation;  properties; 
use,  148;  dependence  of  forma- 
tion on  temperature,  148; 
formed  from  ammonia  by 
use  of  platinized  asbestos 
as  catalyzer,  143,  152;  proper- 
ties; manufacture  by  means 


INDEX 


359 


of  electric  arc,  148;  prepara- 
tion, 153;  presence  after  dyna- 
mite explosion,  228. 

Nitrites,  decomposition  by  acids, 
147,  152. 

Nitrites,    preparation,    147,    152. 

Nitrocellulose,  preparation,  220; 
use,  234. 

Nitrogen,  occurrence  in  air, 
weight  over  square  foot  of 
earth;  in  all  living  bodies, 
source  of,  138,  152;  diagram 
of  course  in  nature;  prepa- 
ration; properties,  139;  oxides 
of,  147;  essential  element  in 
soil,  297. 

Nitrogen  dioxide,  as  catalytic 
agent,  121;  as  oxidizing  agent, 
128;  from  nitric  acid;  pro- 
perties, 144;  formation  "form- 
ula, polymer,  dissociation,  150, 
•  153;  products  on  solution, 
153. 

Nitrogen  oxides,  to  illustrate 
law  of  multiple  proportions, 
50. 

Nitrogen  tetroxide,  reaction  with 
water,  151. 

Nitroglycerine,  manufacture  by 
sulfuric  acid,  125;  manufac- 
ture, how  exploded,  228. 

Nitrous  acid,  preparation,  de- 
composition, 147. 

Nitrous  anhydride,  preparation, 
147,  179;  how  formed,  dissocia- 
tion, 150,  153. 

Nitrous  oxide,  preparation,  prop- 
erties, 147;  use,  153. 

Noble  gases,  properties,  list  of, 
157;  relation  to  halogens  and 
alkali  metals,  158;  zero  group, 
list  of,  160. 


Nomenclature,  binary  com- 
pounds, 15,  acids  and  salts, 
84,  85. 

Non-coking  coals,  properties,  192. 

Normal  salts,  definition,  79, 
118. 

Normal  salts  of  strong  acids  are 
neutral;  of  weak  acids, 
frequently  alkaline,  76. 

Normal  sulfates  of  strong  bases, 
neutral,  125. 

Number  of  molecules,  in  cubic 
centimeter  of  gas;  in  equal 
volumes,  33. 

Oil  of  vitriol,  name  for  sulfuric 

acid,  285. 

Oleic  acid,  in  fats,  227. 
Olein,  derivation,  227. 
Opal,    form    of    silicon    dioxide, 

containing  water,  238. 
Open  hearth  steel,  characteristics 

of,  333;  manufacture,  327. 
Ores  of   common  metals,  reduc- 
tion of,  277. 

Organic   matter,   source   of   car- 
bonic acid,  62. 
Orpiment,  source  of  arsenic  tri- 

sulfide,  179. 
Orthophosphoric    acid,  salts    of; 

tribasic,  174,  183. 
Osmium,  as  catalyser  in  forming 

ammonia,  140. 
-ous,  meaning  of,  84. 
Oxalic    acid,    in    preparation    of 

carbon    monoxide,    211,     216. 
Oxidation,  by  action  of  water  and 

air,   14;  slow,  cause  of  bodily 

temperature,   13;  by  bacteria, 

14. 
Oxidation — reduction,      opposite 

processes,  24. 


360 


INDEX 


Oxides  of  metals,  solution  in 
hydrochloric  acid,  82;  of  anti- 
mony, 181;  of  nitrogen,  list 
of,  147;  of  lead,  list;  properties, 
248;  of  iron,  zinc  and  tin, 
how  reduced,  277. 

Oxidizing  agents,  action  on  hy- 
drochloric acid ;  list  given ;  82, 
90;  potassium  dichromate  as, 
317;  in  preparation  of  potas- 
sium manganate,  321;  use  of 
potassium  permanganate,  322. 

Oxone,  composition,  207. 

Oxy -acetylene  flame,  use  of,  27. 

Oxygen,  preparation  from  red 
oxide  of  mercury,  7 ;  from  potas- 
sium chlorate,  9,  87;  from  liq- 
uid air,  10,  17;  properties,  10, 
17;  occurrence,  12;  velocity  of 
molecules,  34;  weight  of  liter 
of,  43;  per  cent,  in  air,  57;  for- 
mula of,  how  determined,  136; 
in  air,  how  determined,  155, 
160;  preparation  from  barium 
peroxide,  268,  275. 

Oxygen  acids  of  chlorine,  84; 
salts  of,  85. 

Oxyhydrogen  blowpipe,  why  tem- 
perature is  limited,  25;  use  of, 
26. 

Ozone,  preparation,  properties, 
15;  allotropic  form  of  oxygen, 
17;  for  purification  of  water, 
70;  weight  of  liter  of,  137. 

Palmatin,    derivation,    227. 
Palmitic  acid,  in  fats,  227. 
Paper,  manufacture,  220. 
Paraffin,  product  from  petroleum, 

199. 
Paris    green,    composition;   use, 

179,  1S3. 


Parts,  by  weight  and  atoms,  52. 

Peat,  properties,  191,  192. 

Pectin,  preparation ;  use,  224, 234. 

Pectose,  presence  in  fruits,  224, 
234. 

Per-,    meaning    of,    49. 

Perchloric  acid,  properties,  89; 
use  for  detection  of  potassium, 
90. 

"Perfect"  gases,  132. 

Periodic  system,  discussion  of 
tables,  102;  relation  to  atomic 
weights,  133;  tables,  164,  165. 

Periods  of  periodic  system, 
162. 

Perkin,  W.  H.,  discoverer  of 
mauve,  233. 

Permanent  hardness  of  waters, 
263. 

Petrolatum,    see    vaseline,     199. 

Petroleum,  composition;  occur- 
rence, 198;  use;  distillates,  198; 
products  from,  199,  208. 

Phenol,  use  of  benzene  in  manu- 
facture of,  202;  preparation; 
use,  228,  235. 

Phenol  phthalein,  as  an  indicator, 
76. 

Phosphates,  mineral,  use  for 
fertilizers,  125;  use  as  fertilizer, 
170. 

Phosphine,  preparation;  proper- 
ties; comparison  with  ammonia, 
172,  183. 

Phosphonium  iodide,  how  pre- 
pared; comparison  with  am- 
monium iodide,  173. 

Phosphoric  acid,  preparation,  171 ; 
ionization  of,  175;  hydrolysis 
of  salts  of,  175;  by  hydrolysis 
of  phosphorus  pentachloride, 
176. 


INDEX 


361 


Phosphoric  acids,  list  of,  174, 
183. 

Phosphoric  anhydride,  see  phos- 
phorus pentoxide,  173. 

Phosphorous  acid,  by  hydrolysis 
of  phosphorus  trichloride,  176; 
dibasic,  properties,  176,  183. 

Phosphorus,  burning  in  oxygen, 
11;  kindling  temperature  of, 
12;  preparation,  171;  occur- 
rence in  combination,  170; 
properties;  allotropic  forms  of, 
171 ;  in  manufacture  of  matches, 
poisoning  by,  171;  occurrence; 
preparation;  forms,  182;  es- 
sential element  in  soil,  297. 

Phosphorus  chlorides,  hydro- 
lysis of,  176. 

Phosphorus  pentachloride,  prepa- 
ration, 83;  preparation,  hy- 
drolysis, 176,  183. 

Phosphorus  pentoxide,  from  com- 
bustion experiment;  properties, 
11;  gives  acid  with  water,  17; 
preparation;  properties;  use  as 
drying  agent,  173;  preparation; 
use  as  drying  agent,  182. 

Phosphorus  trichloride,  prepara- 
tion, 83. 

Phosphorus  trichloride,  prepara- 
tion, hydrolysis,  176,  183. 

Photographic  plates,  effects  of 
uranium  and  radium  upon, 
268. 

Photography,  use  of  silver  halides 
in,  311. 

Physical    science,    definition,   1. 

Physics,  definition,  1,  5. 

Pig  iron,  manufacture,  324, 
333. 

Plants,  what  growth  depends 
upon,  231.  235. 


Plaster  casts,  264. 

Plaster  of  Paris,  how  made,  110; 
manufacture;  use,  264;  source 
of,  261;  properties,  275. 

Platinic  chloride,  preparation, 
332. 

Platinum,  melting  point,  26; 
as  catalytic  agent,  119;  action 
of*  aqua  regia  upon,  146; 
use;  price;  as  catalyst;  aqua 
regia  as  solvent  of,  331,  332; 
occurrence;  use,  334. 

Plumbic  acid,  salt  of,  248. 

Polymer,    definition,    119. 

Porcelain,  257. 

"Positive"  in  photography  ex- 
planation, 312. 

Positive  and  negative  valences, 
104,  106. 

Potash  salts,  where  found,  68; 
occurrence;  use,  298. 

Potassium,  used  to  decompose 
water,  19;  detection  by  per- 
chloric acid,  90;  solubility  of 
salts  of,  280;  occurrence;  es- 
sential element  in  soil,  297,  307;. 
preparation ;  properties,  298, 
306. 

Potassium  aluminium  sulfate; 
action  in  bread  making,  230. 

Potassium  argenticyanide,  use, 
216,  217. 

Potassium  bicarbonate,  use,  299, 
306. 

Potassium  bromide,  use  in  medi- 
cine, 94. 

Potassium  carbonate,  solution  of, 
alkaline,  213;  in  preparation  of 
potassium  ferrocyanide,  215; 
from  wood  ashes  by  leaching;, 
use  in  soap  making,  297; 
properties,  299;  preparation, 


362 


INDEX 


306;  use  in  preparation  of 
potassium  chromate,  317. 

Potassium  chlorate,  use  for  pre- 
paring oxygen,  9,  87;  chloric 
acid  from,  88;  to  prepare  potas- 
sium perchlorate,  89;  products 
when  heated,  90;  use  in  flash 
light  powders,  283;  prepara- 
tion, use,  299,  306. 

Potassium  chloride,  in  manufac- 
ture of  gunpowder,  279. 

Potassium  chloroplatinate,  prepa- 
ration; use  as  laboratory  test, 
332. 

Potassium  chromate,  preparation, 
properties,  317,  319. 

Potassium  cyanide,  preparation; 
use  in  extracting  gold,  215,  217; 
in  preparation  of  potassium 
ferrocyanide,  215. 

Potassium  dichromate,  to  oxidize 
hydrochloric  acid,  83;  use  in 
manufacture  of  chrome  yellow, 
250;  preparation,  properties, 
use,  317;  use  in  tanning,  318, 
319. 

Potassium  f  erricyanide,  prepara- 
tion, use,  216. 

Potassium  ferrocyanide;  prepa- 
ration, properties,  215,  217; 
electrolysis  of,  216. 

Potassium  fluoride,  use  in  pre- 
paring fluorine,  96. 

Potassium  hydroxide,  by  de- 
composition of  water,  26;  in 
reversible  reaction,  86;  prepa- 
ration, 297,  299;  use  in  prepara- 
tion of  potassium  chromate, 
317. 

Potassium  hypochlorite ;  products 
when  warmed,  90;  how  formed, 
299. 


Potassium  iodide,  use  of,  95,  98. 

Potassium  manganate,  prepara- 
tion, 321. 

Potassium  nitrate,  in  soil,  source 
of,  138;  from  decay  of  organic 
matter;  from  sodium  nitrate, 
144;  preparation  for  manufac- 
ture of  gunpowder,  279;  occur- 
rence; preparation;  use;  299, 
306. 

Potassium  oxide,  preparation, 
298. 

Potassium  perchlorate,  from  pot- 
assium chlorate ;  properties, 
89,  90. 

Potassium  permanganate,  to 
oxidize  hydrochloric  acid,  83; 
preparation,  321;  use,  322. 

Potassium  sodium  tartrate,  prod- 
uct of  action  of  baking  pow- 
ders, 230. 

Potassium  sulfate,  use  in  manu- 
facturing alum,  257. 

Potassium  tartrate,  acid,  229. 

Potential,  difference  between 
metals  and  solutions,  242; 
difference  between  metals,  243; 
difference  in  storage  batteries, 
249. 

Potential  energy,  2. 

Pound,  grams  in  one,  57. 

Precipitates,  formation  of,  rever- 
sible, 113,  127. 

Prefixes,  in  chemical  nomencla- 
ture, 84,  85,  90;  hypo-,  85;  per-, 
49,  85;thio-,  meaning,  126,  296; 
to  distinguish  salts  of  ortho- 
phosphoric  acid,  174. 

Pressure  of  air,  at  sea  level,  36; 
standard  conditions  of,  37. 

Principle  of  van't  Hoff-Le  Chat- 
elier,  120. 


INDEX 


363 


Priestly,  preparation  of  oxygen 
gas  by,  7. 

Producer  gas,  manufacture;  use; 
composition,  204,  208. 

Proteins,  definition;  occurrence, 
230;  effect  of  digestion  upon, 
231. 

Prussian  blue,  use  in  preparation 
of  chrome  green,  319. 

Prussic  acid,  preparation ;  poison- 
ous quality,  215,  217. 

Puddling  process,  for  making 
wrought  iron,  326,  333. 

Pure  substances,  definition ;  3,  5. 

Purification,  by  crystallization 
and  distillation,  2. 

Pyrite,  110. 

Pyroboric  acid,  as  source  of 
borax,  253. 

Pyrophosphoric  acid,  174;  prepa- 
ration of  salts  of,  175. 

Qualitative  analysis,  explana- 
tion; separation  of  metals 
with  insoluble  chlorides,  335; 
groups  of,  112;  separation  of 
metals  with  sulfides  insoluble 
in  dilute  acids,  336;  detec- 
tion of  sulfates,  chlorides,  and 
nitrates,  337. 

Quantitative  analysis,  explana- 
tion, method  of  procedure, 
337.  - 

Quartz,  varieties  of,  238. 

Quinine,  source;  use,  233. 

Radioactive  elements,  disinte- 
gration of  atoms  of,  163; 
detected  by  gold  leaf  electro- 
scope, 269;  estimate  of  life 
of,  271;  three  series  of;  half- 
life  period  of,  275. 


Radiolite  watch  dials,  cause  of 
phosphorescence,  269;  meso- 
thorium  for,  273. 

Radium,  products  of  decomposi- 
tion of,  158;  discovery;  prop- 
erties, occurrence,  268;  price 
of,  269;  products  of,  272;  table 
of  derivatives,  272 ;  how  formed ; 
disintegration  of,  275. 

Radium  sulfate,  properties,  260. 

Raleigh,  discovery  of  argon,  157. 

Ramsay,  discovery  of  argon,  157; 
of  helium  and  niton,  158. 

Ramsay  and  Soddy,  kelium  from 
radium,  272. 

Reactions,  explanation  of  rever- 
sible, 24,  80;  effect  of  removing 
one  product  upon,  115;  endo- 
thermic,  explanation,  201. 

Reduction,  definition,  24;  of 
copper  oxide,  24,  47;  of  ores  of 
common  metals,  277. 

Red  lead,  preparation;  use;  248. 

Red  phosphorus,  properties,  use, 
171,  182. 

Remedies,  for  burns  by  sulfuric 
acid,  124. 

Respiration  calorimeter,  use  of ,  13. 

Reversible  reactions,  explana- 
tion, 24,  80;  how  expressed, 
53;  effect  of  removing  one- 
of  the  products  of,  85;  in 
formation  of  precipitates;  in 
formation  of  ions,  114;  in  pre- 
cipitation of  lead  sulfide,  114;. 
effect  of  temperature  upon, 
120. 

Rochelle  salt,  compositions; 
use  in  preparing  cuprous  oxide, 
309. 

Rock  crystal,  form  of  silicon 
dioxide,  238. 


364 


INDEX 


Roll  brimstone,  definition,  110. 

Rose  quartz,  form  of  silicon 
dioxide,  238. 

Rubies,     composition,     255. 

Rust,  nature;  cause  of,  284. 

Rutherford,  disintegration  of  ele- 
ments, 272. 

Safety  lamp,  197. 

Safety    plugs,    composition,  248. 

Sal    soda,    preparation,  291. 

Saleratus,  see  Potassium  bi- 
carbonate, 299. 

Salt,  how  deposited,  68;  composi- 
tion, occurrence,  72;  formula, 
how  determined;  definition,  75; 
see  sodium  chloride,  290. 

Saltpeter,  source  of  nitric  acid, 
143;  preparation  for  manufac- 
ture of  gunpowder,  279;  source 
of,  291;  occurrence;  prepa- 
ration, use,  299. 

Salts,  definition,  21;  acid,  how 
formed;  normal,  neutral,  and 
acid,  defined,  76,  77;  normal 
of  strong  acids,  are  neutral; 
of  weak  acids,  frequently  aka- 
line  76;  why  called  normal; 
why  called  acid,  78;  neutral; 
alkaline  or  acid,  79;  from  oxy- 
gen acids  of  chlorine,  85;  o 
sulfurous  acid,  preparation; 
properties,  117;  normal,  defini- 
tion, 118;  of  orthophosphoric 
acid,  174;  of  phosphoric  acid, 
hydrolysis  of,  175;  ammonium, 
preparation,  191;  of  metals, 
how  prepared,  reversible  re- 
actions involved,  explanation, 
278;  rules  for  solubility  of,  280; 
reactions  for  preparation  of, 
conditions  necessary ;  when  in- 


soluble, 281;  of  alkali  metals, 
properties,  290;  list  of  am- 
monium, 302;  of  copper,  309; 
of  iron,  329. 

Sand,  composition,  239. 

Saponification,   defined,  227. 

Sapphire,  form  of  corundum,  255. 

Saturated  compounds,  definition, 
199. 

Saturated  solutions,  65. 

Scale,  in  utensils,  cause  of,  62; 
in  boilers,  cause  of,  264. 

Science,  definition,  1,  5. 

Sedimentary  rocks,  composition 
of,  formation,  255. 

Selenite,  properties,  261. 

Selenium,  properties ;  com- 
pounds, use,  126,  128. 

Shales,  composition  of,  forma- 
tion, 255. 

Sicily,  sulfur  from,  109. 

Siderite,     ore    of    iron,     324. 

Silica,  in  manufacture  of  phos- 
phorus, 171. 

Silica,    occurrence   in    soil,    237. 

Silicates,  occurrence  in  soil  and 
rocks,  237;  occurrence;  deriva- 
tion; list  of;  kinds  in  glass,  239; 
list  of,  251;  of  aluminium,  etc., 
rocks  composed  of,  254. 

Silicic  acid,  from  soluble  glass, 
240. 

Silicon,  occurrence  in  crust  of 
earth,  237,  238;  preparation; 
properties;  use,  238;  occur- 
rence, 251. 

Silicon    dioxide,   forms   of,    251. 

Silver,  occurrence ;  separation ; 
use;  properties,  310;  percen- 
tage in  U.  S.  coins,  311;  sepa- 
ration, properties,  314. 

Silver  arsenate,  179. 


INDEX 


365 


Silver  arsenite,  salt  of  arsenious 
acid,  179. 

Silver  bromide,  use  of,  94. 

Silver  bromide,  properties ;  prepa- 
ration; use  in  photography, 
311,  314. 

Silver  chloride,  preparation ; 
properties,  311. 

Silver  iodide,  properties;  prepa- 
ration, 311. 

Silver  nitrate,  preparation; 
properties,  use,  311,  314. 

Slag  of  blast  furnace,  composi- 
tion,'325. 

Slaked  lime,  in  soda  lime,  197; 
preparation;  use,  262;  use  in 
Solvay  process,  293;  use  in 
manufacture  of  sodium  hy- 
droxide, 294. 

Slow  oxidation,  bodily  tempera- 
ture maintained  by,  13. 

Small    calorie,    definition,    58. 

Smokeless  powder,  from  gun 
cotton,  125;  basis  of,  220. 

Smoky  quartz,  form  of  silicon 
dioxide,  238. 

Soap,  manufacture,  saponifica- 
tion,  227;  forms  emulsion  with 
water,  227;  forms  emulsion 
with  water,  235;  action  in 
"hard"  water,  263. 

Soap,  soft,  297;  soft,  hard,  manu- 
facture, 297 ;  soft,  manufacture, 
306. 

Soapstone,  a  silicate,  239. 

Soda  lime,  composition;  use,  197. 

Soddy  and  Ramsay,  helium  from 
radium,  272. 

Sodium,  used  to  decompose 
water,  19;  properties,  73;  solu- 
ble glass  a  silicate  of,  239; 
preparation  by  electrolysis, 


278;  solubility  of  salts  of, 
280;  occurrence;  290,  305; 
sodium  bicarbonate  from;  di- 
sodium  phosphate  from,  291. 

Sodium  acetate,  in  preparation  of 
methane,  197. 

Sodium  aluminate,  preparation, 
256. 

Sodium  ammonium  phosphate, 
175. 

Sodium  antimonite,  preparation, 
181. 

Sodium  bicarbonate,  for  wounds 
by  sulfuric  acid,  124;  use  with 
calomel,  287;  product  in  diges- 
tion; function  in  body,  291; 
preparation  by  Solvay  process, 
293. 

Sodium  bromide,  occurrence ;  elec- 
trolysis, 93;  use  in  medicine,  94. 

Sodium  carbonate,  solution  of, 
alkaline,  213;  to  remove  per- 
manent hardness  in  water, 
264;  source  of,  291;  prepara- 
tion, 291;  product  of  Le  Blanc 
soda  process;  alkalinity  of, 
292;  use  in  manufacture  of 
sodium  hydroxide,  294;  why 
hydrolyzed,  305;  preparation, 
305. 

Sodium  chloride,  electrolysis  of 
fused,  of  aqueous  solution,  72; 
occurrence;  office  in  digestion; 
importance  in  chemical  indus- 
tries, 290,  305. 

Sodium  cyanide,  preparation;  use 
in  extracting  gold,  215,  217. 

Sodium  hydroxide,  by  decompo- 
sition of  water;  affects  red 
litmus,  20,  26;  by  electrolysis,. 
73;  alkaline  reaction,  73; 
formed  by  ionization  of  phos- 


366 


INDEX 


phoric  acid,  175;  in  soda  lime, 
197;  solvent  for  stannic  hy- 
droxide, 246;  source  of,  291; 
preparation;  use,  294,  305; 
properties,  295. 

Sodium  hydroxide,  preparation, 
305. 

Sodium  "hyposulfite,"  use  of 
term,  126;  see  sodium  thio- 
sulfate,  296. 

Sodium  iodide,  occurrence,  elec- 
trolysis, 93. 

Sodium  metaphosphate,  prepara- 
tion, 175. 

Sodium  nitrate,  use  in  manufac- 
ture of  sulfuric  acid,  123; 
occurrence,  use,  144;  in  manu- 
facture of  gunpowder,  279; 
occurrence,  use,  291,  305. 

Sodium  oxide,  preparation,  295. 

Sodium  peroxide,  preparation; 
use,  296,  305. 

Sodium  silicate,  preparation ;  use, 
296. 

Sodium  stannate,  preparation ; 
use,  246. 

Sodium  sulfate,  in  Le  Blanc 
soda  process,  292. 

Sodium  sulfide,  in  Le  Blanc 
soda  process,  292. 

Sodium  sulfite,  preparation,  117; 
preparation;  use,  296,  305. 

Sodium  sulfite,  acid,  preparation, 
116;  use,  117. 

Sodium  thiosulfate,  how  formed, 
properties,  use,  125,  128;  prepa- 
ration, why  so  named,  use, 
296;  use  to  dissolve  salts  of 
silver,  311,  312, 

Soft  soap,  297. 

Soils,  essential  elements  in, 
297. 


Solder,  composition,  245;  alloy 
of  lead,  248. 

Solubility,  in  hot  and  cold  water, 
65;  of  salts,  general  rules  for, 
280. 

Soluble  glass,  preparation;  prop- 
erties, use,  240,  251;  prepara- 
tion; use,  296,  305. 

Solution  pressure  of  metals,  ex- 
planation; cause  of  electric 
current  between  metals,  242, 
251. 

Solutions,  change  of  freezing  and 
boiling  points,  64;  chemical 
activity  in;  ionization  in,  65; 
supersaturated;  saturated,  65, 
69;  neutral  defined,  75,  79. 

Solvay  process  for  preparation 
of  sodium  bicarbonate,  293. 

Specific  gravity  of  elements,  rela- 
tion to  position  in  periodic 
system,  167,  168. 

Spectra,    how    obtained,    302. 

Spectroscope,    explanation,    302. 

Spectrum    analysis,  302. 

Spectrum,  solar,  why  lines  are 
dark,  303. 

Sphalerite,  110;  composition,  284; 
zinc  sulfide,  285;  mineral  source 
of  zinc,  288. 

Spiegeleisen,  use  in  manufac- 
ture of  steel;  preparation, 
321. 

Spontaneous  combustion,  how 
caused;  how  avoided,  14. 

Stannic  acid,  in  fireproofing  cot- 
ton goods,  246;  see  stannic 
hydroxide,  246. 

Stannic  chloride,  preparation; 
properties;  use,  245. 

Stannic  hydroxide,  preparation; 
properties,  use,  246,  251. 


INDEX 


367 


Stannic  oxide,  in  fireproofing 
cotton  goods,  246. 

Stannous  chloride,  preparation; 
properties;  use  as  reducing 
agent,  245. 

Starch,  occurrence;  description; 
use;  granules  of,  221;  test  for 
iodine,  222;  glocuse  from,  223; 
occurrence,  234. 

Stearic  acid,  in  fats,  227. 

Stearin,  derivation,  227. 

Steel,  characteristics  of;  Bes- 
semer; open  hearth,  326;  what 
hardness  depends  upon,  333. 

Steel  furnace,  open  hearth,  tem- 
perature of,  26. 

Stereotype  metal,  use  of  bismuth 
in,  181. 

Stibine,  preparation;  properties, 
180,  184. 

Stibnite,  source  of  antimony; 
properties,  179. 

Still,  to  concentrate  alcohol,  225. 

Storage  batteries,  construction, 
248;  action  of;  charging  of; 
difference  in  potential  between 
plates,  248;  exhaustion,  how 
tested,  249;  substances  formed 
in,  252. 

Storage  of  gases,  9. 

Stove   polish,   graphite,  187. 

Strontium,  occurrence;  proper- 
ties, 268. 

Strychnine,  properties,  233. 

Sublimation,  definition,  94. 

"Subnitrate"  of  bismuth,  see 
bismuth,  basic  nitrate,  182. 

Substances,  pure,  definition,  2; 
not  volatile,  change  in  freezing 
and  boiling  points  by,  64. 

Sucrose  or  cane  sugar,  prepara- 
tion, 222,  234. 


Suffixes,  in  chemical  nomencla- 
ture, 90;  -ate,  85;  -ide,  15; 
-ic,  84;  -ite,  85;  -ous,  84. 

Sugar,  kinds  of;  preparation,  222. 

Sugar  of  lead,  see  lead  acetate, 
250. 

Sulfantimonite,  ammonium,  181. 

Sulfarsenite,     ammonium,     179. 

Sulfates,  list  of  insoluble,  125, 
128. 

Sulfides  in  analysis,  112;  pre- 
cipitation due  to  ions,  113; 
effect  of  strength  of  acids  in 
precipitation  of,  115;  of  zinc 
and  antimony,  how  reduced, 
277. 

Sulfides  of  copper  and  lead,  how 
reduced,  278. 

Sulfur,  burning:  in  oxygen,  11; 
occurrence,,  how  secured,  109; 
occurrence  as  sulfates  and 
sulfides  r  110';  use  in  sulfuring 
fruit;  in  lime-sulfur  wash;  in 
gunpowder;  in  mamuf  acture ; 
forms  of r  110;, use  in  dyes,  111; 
allotropie-  forms,.  111  ;liquid  and 
gaseousr  propertiesy  111;  boil- 
ing point,.  112;  occurrence;  use; 
allotropie  forrnsr  127. 

Sulfur  dioxide^  by  burning  sul- 
fur in  oxygen;  properties,  11; 
gives  acid  with  water,  17; 
use  in  sulfuring  fruit,  '110; 
occurrence ;  properties ;  prepara- 
tion, 115;  laboratory  prepara- 
tion ofr  118;  preparation  from 
iron  pyrites,  122;  preparation; 
user  127;  preparation,  145. 

Sulfur  trioxide,  formation  from 
sulfur  and  oxygen  reversible; 
preparation;  properties,  118, 
119r 


368 


INDEX 


Sulfuric  acid,  with  zinc;  with  iron, 
21;  electrolysis  of,  42;  use  to 
liberate  chlorine,  87 ;  from  iron 
pyrites,  110;  manufacture  by 
contact  process,  119;  prepara- 
tion by  lead  chamber  process, 
121. 

Sulfuric  acid,  concentration  of; 
properties,  123;  charring  by; 
wounds  by,  remedies;  use  of, 
124;  use  in  preparation  of 
nitric  acid,  144;  as  a  catalyzer, 
211;  behavior  in  storage  bat- 
teries, 249;  use  in  Le  Blanc 
soda  process,  292. 

Sulfuric  anhydride,  see  sulfur 
trioxide,  120. 

Sulfuring  fruit,  with  sulfur  diox- 
ide, 115. 

Sulfurous  acid,  preparation, 
1 17 ;  dibasic ;  salts  of,  117;  prepa- 
ration; properties,  128. 

"Super-phosphate"  of  calcium, 
composition,  265 ;  use  as  fer- 
tilizer, 275. 

Supersaturated   solutions,    65. 

Symbols    of    elements,    51,    56. 

Sympathetic  ink,  use  of  cobalt 
chloride  in,  331,  334. 

Synthesis — analysis,  explanation, 
42,  55. 

Synthesis,  definition,  3;  of  com- 
pounds, 5;  of  ammonia,  142; 
of  nitric  oxide,  148. 

Talc,  a  silicate,  239. 

Tan  bark,  use  in  manufacture  of 
white  lead,  250. 

Tannin,  source;  use  in  tanning, 
318. 

Tanning,  chrome,  use  of  potas- 
sium dichromate  ;n,  317,  319. 


Tar  distillates,  source;  use,  189. 

Tartaric  acid,  preparation ;  source ; 
salts,  228,  235. 

Tellurium,  properties,  occurrence, 
126,  128;  exception  in  periodic 
table,  168. 

Temperature,  lowest  obtained, 
35;  how  measured,  34,  40; 
absolute,  explanation,  34,  40; 
standard  conditions  of,  37; 
influence  on  rate  of  combina- 
tion of  elements  in  reversible 
reaction,  120. 

Temporary  hardness  of  waters, 
263. 

Tetraphosphorus  trisulfide,  use 
in  the  manufacture  of  matches, 
172,  182. 

Thermite.process,  Goldschmidt's, 
256,  258. 

Thermometers,  graduation  of, 
34;  use  of  mercury  in,  286. 

Thio-,  meaning,   126,  296. 

Thorium,  series  of  derivatives, 273. 

Thyroid    gland,    iodine    in,    95. 

Tin,  occurrence;  preparation; 
properties; use,  241,  251 ;  alloys 
of,  description,  245;  alloys  of, 
list  of;  chlorides  of,  251;  re- 
duction of  oxides  of,  277. 

Tinware,    cause    of    rust,     243. 

Toluene,  derivative  of  coal  tarr 
189;  use  in  making  trinitrotol- 
uene (T.  N.  T.)  202. 

Toxins,  how  formed;  properties, 
232,  235. 

Tri-,   prefix,  174. 

Tribasic  acid,   defined,   76,    118. 

Tricalcium  phosphate ;  occur- 
rence; 170. 

Trinitrotoluene  (T.  N.  T.),  prepa- 
ration; u^e,,  202. 


INDEX 


369 


Trisilver  arsenate,  179. 
Tungsten,    why    suggested    for 

electric   lights,    168;  discovery 

of  use,  318,  319. 

Type-metal,  composition  of,  180. 
Typhoid  fever,  from  impure  water, 

64. 

Ultra-violet  light  for  purification 
of  water,  70. 

Unit  volume,  measure  of,  132. 

Unit  volume  of  gases,  how  deter- 
mined, 132. 

Unsaturated  compounds,  defini- 
tion, 199. 

Uranium,  as  catalyser  in  forming 
ammonia,  140;  effect  of  com- 
pounds on  photographic  plates, 
268. 

Valence,  negative  in  halogens,  99; 
explanation,  101;  nomencla- 
ture of,  102;  varying,  103; 
positive  and  negative,  104, 
107;  table  of,  106;  table  of 
common,  107;  definition;  ex- 
pressed by  graphical  formulas, 
107;  indicated  by  position  of 
element  in  periodic  system, 
163. 

Van't  Hoff-Le  Chatelier,  principle 
of,  120,  128;  application  of 
principle  of,  120,  142,  148. 

Vapor  pressure  of  water,  60; 
explanation,  62;  law  of;  table 
for  ice  and  water,  63. 

Vaporization  of  water,  latent  heat 
of,  59,  69. 

Vaseline,  product  from  petroleum, 
199. 

Velocity  of  molecules,  law  of,  -33; 
of  hydrogen;  of  oxygen,  33, 

24 


Ventilation,   standard   of,  159. 

Vinegar,  product  of  acetic  fer- 
mentation, 227,  235;  method 
of  manufacture,  227. 

Vitriols,compounds  so  named,285. 

Volatile  compounds,  conditions 
for  formation  of,  278. 

Volume  of  gas,  variation  with 
temperature,  35;  reduction  of 
gases  to  standard,  37;  gram 
molecular  weight  of;  measure 
of,  134. 

Volumes  of  gases,  number  of 
molecular  in  same,  33,  40; 
law  of  change  of;  method  of 
finding,  37,  40;  ratio  in  simple 
numbers,  132. 

Washing  soda,  see  sodium  car- 
bonate, 291. 

Water,  decomposition  of;  by  red 
hot  iron,  19,  26;  by  sodium  or 
potassium,  19;  by  electricity, 
44. 

Water,  composition  by  weight, 
44,  45,  50;  volumetric  compo- 
sition of,  44;  supercooled,  58; 
boiling  of;  superheated;  latent 
heat  of  vaporization  of,  59; 
expansion  on  freezing,  59; 
cooling  by  convection;  vapor 
pressure  of,  60,  69;  natural,  61; 
"hard,"  62;  table  of  vapor  pres- 
sures, 63;  as  disease  carrier, 
64,  69;  mineral,  67;  latent  heat 
of  vaporization  of;  maximum 
density,  69;  ionization,  74; 
methods  of  purification,  70; 
hard,  explanation,  263;  tem- 
porary and  permanent  hard- 
ness, explanation;  how  re- 
moved, 264. 


Water  gas,  preparation;  proper- 
ties; poisonous  quality,  203r 
208. 

Water  vapor  in  air,  156. 

Weight,  no  change  of  in  burning, 
7. 

White  lead,  preparation;  prop- 
erties; use,  250.;  composition; 
use,  252. 

White  vitriol,  name  for  zinc 
sulfate,  285. 

Window  glass,  composition,  239. 

Wire  gauze,  use  in  safety  lamp, 
197. 

Wood  material,  changes  during 
geological  time,  192. 

Wood's  metal,  composition,  use, 
181. 

Wrought  iron,  manufacture;  use, 
326,  333. 

Xenon,  where  found,  158. 
X-rays,  effects  similar  to  those  of 

radium,    269. 
X-ray      spectra,     to     determine 

atomic  number,,  168. 

Yeast,  cause  of  fermentation,  224. 


Zero,,  absolute,.  explanation,. 
34. 

Zero  grotrpr  noble  gases,  list  of, 
160. 

Zinc,  for  preparation  of  hydrogen, 
21;  as  eoating  for  iron,  243; 
reduction  of  oxides  of;  reduc- 
tion of  sulfides  of,  277;  occur- 
rence; preparation;  properties; 
use;  alloys  of,  284,  288;  use  in 
batteries,  285;  use  in  refining 
silver,  310;  use  in  cyan-'de 
process  for  gold,  312;  use  in 
nickel  silver,  331. 

Zinc  chloride,  how  formed,  21; 
how  prepared,  83;  use  as 
disinfectant,  288. 

Zinc  hydroxide,  properties,  260. 

Zinc  oxide,  preparation;  prop- 
erties; use;  when  preferable 
to  white  lead  in  paint,  285, 
288. 

Zinc  sulfate,  how  formed,  21; 
popular  name  of,  288. 

Zinc  sulfide,  action  of  radium 
upon,  269;  use  of  phosphores- 
cent screen  of,  274;  occurrence; 
preparation,  285;  color,  285. 


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-JWSf  #  III* 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 
This  book  is  DUE  on  the  last  date  stamped  below. 


...  ... 


due.- 


OCT    14  1947 

24May'49SL 


lMar55BP 


JUN9    '61    R 


LD  21-100m-12.'46(A2012sl6)4120 

1954  til 


IEP1 


VB   1694. 


THE  UNIVERSITY  OF  CALIFORNIA  LIBRARY 


